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ELEMENTARY  CHEMISTRY 


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HENEI  MOISSAN 
B.  Paris,  Sept.  28,  1852. 


(See  pages  27G,  319 


ELEMENTARY 
CHEMISTRY- 


BY 

ROBERT  HART  BRADBURY,  A.M.,  PH.D. 

TEACHER  OF  CHEMISTRY,  CENTRAL  MANUAL  TRAINING-SCHOOL, 

PHILADELPHIA;   FORMERLY    LECTURER   ON   PHYSICAL 

CHEMISTRY,  DEPARTMENT  OF  PHILOSOPHY, 

UNIVERSITY  OF  PENNSYLVANIA 


NEW    YORK 

D.    APPLETON     AND     COMPANY 
1909 


COPYRIGHT,  1903 
BY   D.    APPLETON   AND   COMPANY 


A- 


PKEFACE 


IN  this  book  I  have  attempted  to  write  a  brief  account 
of  the  present  state  of  chemical  science  for  the  use  of  those 
beginning  the  subject  in  colleges  and  secondary  schools.  I 
have  tried  to  make  the  book  represent  the  subject  as  it  is 
to-day,  and  I  believe  that  everything  that  lies  within  its 
scope  in  the  later  literature  has  been  considered  with  a  view 
to  its  admission  or  exclusion.  It  is  difficult  to  avoid  errors 
in  a  work  of  this  kind,  and  I  can  hardly  hope  that  I  have 
completely  succeeded  in  doing  so,  but,  at  all  events,  any 
blunders  which  the  book  may  contain  are  my  own  property, 
for  I  have  not  consciously  copied  after  any  existing  manual, 
and  I  have  settled  all  cases  of  doubt  as  to  the  correctness 
of  a  statement  by  critical  consideration  of  the  original 
memoir. 

The  book  may  fairly  claim  to  be  regarded  as  an  inde- 
pendent account  of  the  subject.  It  is  the  outgrowth  of 
about  ten  years  of  high-school  teaching,  and  I  have  been 
careful  to  include  nothing  in  it  which  I  have  not  taken 
up  successfully  and  repeatedly  with  ordinary  classes.  This 
is  the  only  safe  criterion,  and  I  have  felt  compelled  to  fol- 
low it,  though  it  has  resulted  in  the  exclusion  of  some 
generalizations  which  I  should  have  been  glad  to  admit. 
I  believe  that  in  this  way  I  have  entirely  avoided  the  dan- 
ger of  becoming  unintelligible  to  the  beginner. 


viii  ELEMENTARY  CHEMISTRY 

pleasure  to  express  here  my  gratitude  for  the  care  with 
which  he  has  carried  out  this  exacting  task. 

Prof.  Edgar  F.  Smith,  director  of  the  Chemical  Labora- 
tory of  the  University  of  Pennsylvania,  has  been  good 
enough  to  read  the  manuscript  and  to  favor  me  with  his 
opinions  as  to  its  arrangement  and  availability  as  a  text- 
book for  beginners. 

Dr.  E.  A.  Partridge,  of  the  Central  Manual  Training 
School,  has  read  the  work  both  in  proof  and  manuscript  and 
has  made  a  number  of  valuable  suggestions. 

The  Laboratory  Manual  aims  to  supply  a  serviceable 
outline  of  experimental  work  to  accompany  the  text.  The 
precise  directions  which  are  given  for  the  performance  of 
the  experiments  will,  I  trust,  save  the  teacher  much  useless 
repetition,  while  the  constant  precautionary  suggestions 
will  somewhat  lighten  his  burden  of  responsibility.  If  the 
directions  are  properly  followed,  the  experiments  can  be 
carried  out  without  danger  of  any  kind. 

I  believe  it  to  be  unwise  to  place  phosphorus  in  the 
hands  of  beginners,  and  the  substance  is  employed  only  once 
in  the  course.  In  this  instance  the  danger  is  slight,  but 
the  experiment  can  be  carried  out  on  the  lecture-table  if 
preferred. 

Almost  all  the  experiments  can  be  performed  with 
very  simple  and  inexpensive  materials.  Perhaps  half  a 
dozen  require  somewhat  expensive  apparatus — the  Hoff- 
mann apparatus  for  electrolysis,  a  U-shaped  endiometer  with 
three-way  stop-cock,  an  induction  coil  and  two  burettes  are 
the  most  important.  Some  will  prefer  to  carry  out  these 
experiments  on  the  lecture-table.  However,  I  have  not 
labeled  them  "  teacher's  experiments,"  since  I  believe  the 
decision  should  be  left  to  the  teacher  himself. 

The  object  of  the  questions  which  are  scattered  through 
the  manual  is  to  afford  assistance  in  the  difficult  and  essen- 
tial task  of  making  the  student  do  his  laboratory  work  in 


PREFACE  IX 

a  thoughtful  spirit.  The  materials  for  answering  the  ques- 
tions are  all  contained  either  in  the  text  or  in  the  laboratory 
exercises,  and  they  should  be  answered  by  the  student,  not 
by  the  teacher,  save  as  a  last  resort. 

In  the  preparation  of  the  Laboratory  Manual  free  use 
has  been  made  of  the  available  publications  on  the  subject 
of  experimental  chemistry.  The  recommendations  of  the 
College  Entrance  Examination  Board,  the  Committee  of 
Ten  on  Secondary  School  Studies,  of  the  Board  of  Eegents, 
and  of  Professor  Richards,  of  Harvard,  have  been  carefully 
considered  and  have  afforded  me  much  valuable  assistance. 
About  thirty  of  the  cuts  are  after  drawings  made  by  my 
wife,  and  six  from  original  photographs ;  the  remainder  are 
taken  from  various  sources. 

R.  H.  B. 

CENTRAL  MANUAL  TRAINING  SCHOOL,  PHILADELPHIA. 


CONTENTS 


CHAPTER  PAGE 

I.— WATER .1 

II.— SOLUTION 12 

III. — PHYSICAL  AND  CHEMICAL  CHANGE 23 

IV. — MIXTURE — ELEMENT — COMPOUND 31 

V. — HYDROGEN .35 

VI. — OXYGEN  AND  HYDROGEN  PEROXIDE 40 

VII. — COMBUSTION 51 

VIII. — NAMING  CHEMICAL  COMPOUNDS — CHEMICAL  SYMBOLS  AND 

EQUATIONS — DIFFERENT  KINDS  OF  CHEMICAL  CHANGE   .  58 

IX. — SALT  AND  SODIUM 64 

X. — CHLORINE    . 75 

XI. — THE  CHLORIDES — COMPOUNDS  OF  CHLORINE  CONTAINING 

OXYGEN      .           . 84 

XII. — THE    ATOMIC    THEORY — THE    LAW    OF    MULTIPLE     PROPOR- 
TIONS            •      .  90 

XIII. — THE   ATMOSPHERE — NlTROGEN 98 

XIV. — COMPOUNDS  OF  NITROGEN  AND  HYDROGEN         .        .        .  105 

XV. — COMPOUNDS  CONTAINING  NITROGEN  AND  OXYGEN       .        .111 

XVI. — ATOMIC  AND  MOLECULAR  WEIGHTS — AVOGADRO'S  RULE    .  119 

XVII. — ACIDS,    BASES,    AND     SALTS — ELECTROLYTIC     DISSOCIATION 

— METALS  AND  NON-MI  TALS 125 

XVIII. — THE  SODIUM  GROUP 135 

XIX. — THE  COPPER  GROUP 144 

XX.— THE  COPPER  GROUP  (cc  .tinned) 152 

XXI. — THE  COPPER  GROUP  (Continued) 159 

XXII.— THE  CALCIUM  GROTTP 163 

XXIII. — THE  ZINC  GROUP 172 

XXIV.— THE  ZINC  GROUP  (continued) 177 

XXV.— MERCURY     .                181 

XXVI. — BORON  AND  ALUMINIUM 187 

xi 


Xll  ELEMENTARY  CHEMISTRY 

CHAPTER  PAGE 

XXVII.— THE  CARBON  GROUP 195 

XXVIII.— THE  CARBON  GROUP  (continued) 201 

XXIX.— LEAD 206 

XXX. — THE  NITROGEN  GROUP 210 

XXXI. — OXIDES  AND  ACIDS  OF  PHOSPHORUS — HALOGEN   COM- 
POUNDS         216 

XXXII. — THE  NITROGEN  GROUP  (continued)       ....  220 
XXXIII.— THE  NITROGEN  GROUP  (continued)       .        .        .       .225 

XXXIV. — THE  CHROMIUM  GROUP 229 

XXXV. — THE  OXYGEN  GROUP 234 

XXXVI. — MANGANESE 246 

XXXVIL— THE  HALOGENS »      .        .250 

XXXVIII. — BROMINE  AND  IODINE    . 255 

XXXIX.— IRON 261 

XL. — COBALT  AND  NICKEL — THE  PLATINUM  METALS    .        .  268 

XLL— CARBON 272 

XLII. — CARBON  DIOXIDE — CHEMICAL   ENERGY — CARBON  MON- 
OXIDE— CARBON  DISULPHIDE 278 

XLIII. — SOME  CARBON  COMPOUNDS 286 

XLIV. — SOME    ADDITIONAL    CARBON     COMPOUNDS — CHEMICAL 

PROCESSES   OF   THE   ANIMAL  BODY       ....  297 

XLV.— THE  PERIODIC  LAW 303 

XL VI. — THE  HISTORY  OF  CHEMISTRY 311 

INDEX                                                                           ,  321 


LIST    OF    PORTRAITS 


FACING 
PAGE 

Henri  Moissan Frontispiece 

Joseph  Priestley 30 

Sir  Humphry  Davy 68 

Daniel  Rutherford 98 

Antoine  Laurent  Lavoisier 126 

John  Dalton 155 

Friedrich  Wohler 189 

Justus  von  Liebig 227 

Robert  Wilhelm  Bunsen 262 

Dimitri  Ivanovitch  Mendelejeff  .        .        .        .        .        .        .300 

xiii 


ELEMENTARY  CHEMISTRY 


CHAPTER  I 

WATER 

1.  EVERY  one  is  familiar  with  the  bodies  called  ice,  liquid 
water,  and  steam.    Most  people  are  aware,  also,  that  these 
three  things  are  not  regarded  as  different  substances,  but 
merely  as   three   different   forms   of  the  same   substance. 
Now,  a  clear  idea  of  the  meaning  of  this  statement  is  of 
the   first   importance   in   chemistry   as   well   as   in    other 
branches  of  knowledge.    At  the  outset,  therefore,  it  is  de- 
sirable to  make  a  careful  study  of  ice,  water,  and  steam, 
and  then  to  consider  what  we  mean  by  the  statement  that 
they  are  different  forms  of  the  same  thing. 

2.  Occurrence  of  water. — Water  is  abundant  in  nature. 
When  an  excavation  is  made  to  a  sufficient  depth,  water  col- 
lects in  the  lower  portion.     In  this  way  we  become  aware 
that  everywhere  below  a  certain  level — called  the  water-line 
— the  rocks  are  saturated  with  water.    Of  course  the  water- 
line,  the  upper  limit  of  the  water-  in  the  rocks,  is  not  a  line, 
but  a  surface.    Its  depth  varies  much  in  different  places,  and 
at  the  same  place  in  different  seasons.     Where  it  rises  to 
the  surface  of  the  soil  there  is  a  body  of  water,  river,  lake, 
or  ocean,  according  to  the  conditions,  while  in  deserts  and 
dry  regions  generally  it  is  met  with  at  some  distance  be- 
low the  ground.    Many  rocks — e.  g.,  granite — contain  large 

1 


2  ELEMENTARY  CHEMISTRY 

quantities  of  water  in  the  form  of  minute  drops,  enclosed  in 
the  crystals  which  make  up  their  mass. 

Some  rocks — e.  g.,  serpentine — are  compounds  contain- 
ing water.  In  these  the  presence  of  the  water  is  not  evi- 
dent on  inspection,  but  becomes  plain  when  the  water  is 
liberated  in  some  way.  When,  for  example,  a  small  piece 
of  such  a  rock  is  heated  in  a  glass  tube  sealed  at  one  end, 
drops  of  water  condense  in  the  cool  part  of  the  tube.  In 
this  state  of  chemical  combination  water  makes  up  an  im- 
portant part  of  the  mass  of  the  earth's  crust. 

The  water  of  nature  is  never  pure.  Even  rain-water 
dissolves  substances  from  the  atmosphere,  that  which  falls 
at  the  beginning  of  the  rain  being  more  impure  than  the 
later  portions,  and,  of  course,  the  water  that  oozes  through 
the  rocks  takes  up  from  them  (dissolves)  a  great  variety  of 
substances.  The  purest  natural  water  is  that  which  results 
from  the  melting  of  ice. 

3.  Preparation  of  pure  water. — In  order  to  obtain  a 
water  pure  enough  for  the  purposes  of  the  chemical  labora- 


FIG.  1.— Distillation  of  water. 


tory,  natural  water  is  converted  into  steam  by  being  heate( 
in  a  copper  vessel,  B,  and  the  vapor  is  passed  into  a  tube 
of  tin  or  copper,  D,  which  is  kept  cool  by  water  circulating 


WATER  3 

outside  it.  Here  the  steam  condenses,  the  earthy  impurities 
have  been  left  behind  in  the  vessel  in  which  the  water  was 
boiled,  and  the  purer  water  is  collected  in  another  vessel,  0. 
This  process  is  called  distillation,  and  is  frequently  applied 
to  the  purification  of  other  substances  (Fig.  1). 

Distilled  water  contains  traces  of  substances — ammonia 
and  nitrogen,  for  example — from  which  it  is  difficult  to  free 
it.  For  certain  special  purposes  it  can  be  further  purified 
by  a  process  of  fractional  freezing.  The  water  is  partially 
frozen  and  the  liquid  rejected.  Then  the  ice  is  melted  and 
the  partial  freezing  repeated,  and  so  on  for  eight  or  ten 
operations.  In  this  way  water  of  extraordinary  purity  can 
be  obtained,  but  never  perfectly  pure.  In  fact,  such  a 
thing  as  an  "  absolutely  pure  substance  "  has  never  been 
obtained.  All  that  our  methods  of  purification  can  do  is 
to  reduce  the  impurities  to  such  small  quantities  that  we  are 
unable  to  detect  them,  and  they  have  no  effect  on  the  be- 
havior of  the  substance. 

4.  Physical  properties. — Thin  layers  of  water  appear 
transparent  and  colorless,  but  in  a  layer  of  two  meters  or 
over  it  has  a  distinct  blue  color.  A  beam  of  light  passed  into 
water  is  just  as  invisible  as  in  a  vacuum  or  in  dust-free  air. 
At  all  temperatures  water  or  ice  gives  off  into  the  space  in 
contact  with  it  the  vapor  of  water.  If  this  space  is  enclosed, 
if,  for  example,  the  water  is  placed  in  a  dish  and  covered 
with  a  bell- jar,  then  the  water  will  evaporate  at  any  fixed 
temperature  until  every  unit  of  volume  of  the  enclosed 
space  contains  a  definite  quantity  of  the  vapor.  The  higher 
the  temperature  the  more  water-vapor  exists  in  every  cubic 
centimeter  of  the  space,  but  at  any  fixed  temperature  the 
quantity  is  always  the  same.  In  other  words,  the  vapor 
pressure  of  water  in  contact  with  the  liquid  is  always  the 
same  at  the  same  temperature,  and  this  pressure  increases 
as  the  temperature  rises.  This  is  the  form  in  which  the 
statement  is  usually  made,  and  it  is  true  not  only  of  water, 
2 


4  ELEMENTARY  CHEMISTRY 

but  also  of  all  pure  substances,  solid  or  liquid,  which  evapo- 
rate unaltered. 

In  open  vessels  the  water  is  diffused  into  the  surround- 
ing air,  and  no  such  saturation  of  the  space  can  occur; 
hence  the  water  evaporates  completely.  Heated  in  open 
vessels  the  vapor-pressure  of  the  water  rises  until  at  100°  it 
equals  the  average  pressure  of  the  air,  and  the  liquid  boils. 
When  boiling  has  once  begun  the  total  heat  supply  is  con- 
sumed in  converting  the  water  into  steam,  asjjJ  no  rise  in 
temperature  occurs  until  the  liquid  has  disappeared.  This 
boiling  away  at  a  constant  temperature,  pressure  remaining 
the  same,  is  common  to  all  pure  substances.  Mixtures  of 
two  substances — e.  g.,  a  mixture  of  alcohol  and  water — 
often  show  a  gradual  rise  in  the  boiling  temperature  as  the 
quantity  of  liquid  decreases,  but  there  are  many  mixtures 
which  boil  away  at  a  constant  temperature,  like  a  pure 
substance. 

5.  Freezing  of  water. — When  water  at  the  ordinary  pres- 
sure is  cooled  to  0°  it  begins  to  solidify.  No  fall  of  tem- 
perature takes  place  during  the  freezing  until  all  the  water 
has  become  ice.  By  proper  precautions  freezing  can  be 
avoided  and  the  liquid  cooled  to  — 10°  or  lower,  but  as 
soon  as  it  begins  to  freeze  the  temperature  rises  tqjpb0,  and 
when  ice  melts  the  temperature  remains  the  same^intil  the 
ice  is  all  melted.  It  is  impossible  to  heat  ice  above  0°. 

All  pure  substances  behave  like  wateji;  *  each  melts  and 
solidifies  at  a  constant  temperature.  However,  many  mix- 
tures behave  in  the  same  way.  And  all  solids  which  are 
pure  single  substances  are  like  ice  in  this  respect,  that  it  is 
impossible  to  heat  them  above  their  melting-points. 

1  In  this  paragraph  and  the  preceding  one  it  is  assumed  that  the 
substance  is  heated  in  open  vessels — that  is,  under  the  pressure  of  the 
atmosphere.  In  most  cases  the  temperature  at  which  a  solid  melts  is 
raised  when  the  pressure  is  increased.  Ice  is  an  exception.  Its  melt- 
ing temperature  is  lowered  when  the  pressure  is  raised. 


WATER  5 

6.  Action   of  pure   water   on   the   system. — Electrical 
resistance. — The  purest  water  is  a  poison  to  micro-organ- 
isms, and  causes  vomiting  and  other  disturbances  in  the 
human  system.     Distilled  water  is  too  pure  to  be  whole- 
some, and  certain  waters  which  occur  in  nature,  for  instance 
the  water  formed  by  the  melting  of  the  ice  of  the  Swiss 
glaciers,  are  so  pure  that  they  are  quite  unfit  for  use. 

Water  offers  an  enormous  resistance  to  the  passage  of 
the  electric  current.  In  passing  through  a  column  of  water 
one  millimeter  in  length  the  current  experiences  as  much 
resistance  as  in  a  copper  wire  of  the  same  cross-section 
long  enough  to  encircle  the  earth  at  the  equator  a  thou- 
sand times. 

7.  Ice. — Ice,  like  water,  is  blue  in  large  masses.     It  is 
composed  of  small  imperfect  crystals ;  and,  according  to  the 
size  of  these,  the  properties  of  the  mass  vary  slightly.    Thus, 
when  water  has  been  cooled  below  the  freezing-point  and 
then  caused  to  freeze  suddenly  by  agitation,  the  ice  is  whiter, 
softer,   less   transparent,   and  more  filled  with  air-bubbles 
than  ordinary  ice.     The  most  obvious  difference  between  ice 
and  liquid  water  is  that  ice  is  a  solid,  by  which  we  mean  that 
it  has  a  .shape  of  its  own  which  it  retains,  while  water  takes 
the  shapei  of  the  containing  vessel. 

SteairiStn.  thin  layers  appears  colorless.  In  large  masses 
it  is  blue;  and  the  color  of  the  sky  is,  in  part,  due  to  the 
water-vapor  of  th-^'jjfr.  The  white  cloud  commonly  called 
steam  consists  of  fine  drops  of  water.  The  density  of  steam 
is  only  a  little  more  than  one-half  that  of  the  air.  This  low 
density  is  one  important  distinction  between  vapors  and  the 
liquids  from  which  they  are  obtained.  Under  ordinary 
pressure  the  vapor  is  always  immensely  lighter  than  the 
liquid.  Another  difference  is  that  the  liquid  left  to  itself 
has  a  limiting  upper  surface,  the  vapor  has  not. 

8.  Composition  and  chemical  behavior. — There  are  two 
important  methods  of  investigating  the  composition  of  a 


6 


ELEMENTARY  CHEMISTRY 


substance.  Either  it  is  separated  into  its  constituents, 
which  is  analysis,  or  the  constituents  are  brought  together 
and  the  substance  produced.  This  is  called  synthesis.  Both 
methods  can  be  employed  with  water.  Water  can  be  sepa- 
rated into  its  components  by  heat,  but  the  temperature  re- 
quired is  extremely  high.  The 
most  convenient  method  of 
analysis  is  furnished  by  the 
electric  current. 

9.  Electrolysis  of  water. — 
The  apparatus  (Fig.  2)  has  a 
central  tube,  which  serves  as 
a  funnel  for  filling  it  with 
liquid;  in  the  two  side-tubes 
collect  the  gases  which  are 
given  off  during  the  process. 
The  current,  coming  from 
any  suitable  source,  passes  in- 
to platinum  wires  which  are 
fused  through  the  glass  and  end  in  small  pieces  of  platinum 
foil,  one  in  each  of  the  side-tubes.  These  pieces  of  foil  are 
called  the  electrodes,  and  it  is  there  that  the  products  of 
the  decomposition  appear.  It  has  been  said  that  the  elec- 
trical resistance  of  pure  water  is  enormous.  For  this  reason 
the  apparatus  is  filled  with  water,  to  which  there  has  been 
added  about  one-tenth  of  its  volume  of  sulphuric  acid,  which 
greatly  increases  the  conductivity  of  the  liquid.  It  is 
impossible  to  explain  the  role  of  the  sulphuric  acid  in  an 
elementary  work,  further  than  to  say  that  while  it  conducts 
the  current  it  is  found  unaltered  after  the  experiment,  and 
only  the  water  is  decomposed.  As  soon  as  the  current  is 
made  to  pass,  it  is  seen  that  bubbles  of  colorless  gas  collect, 
and  are  evolved  from  both  electrodes.  When  some  gas  has 
collected  the  stop-cocks  at  the  top  may  be  opened  and  a 
burning  match  applied  to  the  issuing  gases.  One  is  com- 


FIG.  2.— Electrolysis  of  water. 


WATER  7 

bustible,  burning  with  a  pale  blue  flame;  the  other  is  not, 
but  if  the  flame  is  extinguished  and  the  match-stick,  st'ill 
bearing  a  spark,  is  held  in  the  gas  the  wood  bursts  into 
flame  at  once.  The  combustible  gas  is  called  hydrogen,  and 
the  one  which  supports  combustion  so  much  better  than  air 
is  oxygen. 

Now  the  stop-cocks  may  be  closed  and  the  gases  allowed 
to  accumulate.  It  will  be  seen  that  twice  as  much  hydrogen 
collects  at  the  negative  electrode  as  of  oxygen  at  the  posi- 
tive. 

Thus,  when  water  is  decomposed  by  the  electric  cur- 
rent, two  volumes  of  hydrogen  and  one  of  oxygen  appear. 

10.  Synthesis  of  water. — The  conclusion  that  water  is 
composed  of  these  two  gases  in  these  proportions  can  be  cor- 


FIQ.  3.— Formation  of  water  by  burning  hydrogen. 

roborated  by  synthesis.  That  water  is  a  compound  of 
hydrogen  and  oxygen  can  be  shown  in  a  simple  manner  by 
burning  a  jet  of  dry  hydrogen  under  a  cold  jar,  when 
drops  of  water  condense  on  the  glass  (Fig.  3).  This  ex- 
periment gives  us  no  information  about  the  quantities  in 
which  the  gases  enter  into  the  compound.  Such  informa- 
tion is  readily  obtained  by  means  of  the  apparatus  (Fig.  4). 
The  curved  tube  is  called  an  eudiometer.  It  is  first  filled 


8 


ELEMENTARY   CHEMISTRY 


FIG.  4.— Synthesis  of  water. 


with  mercury,  and  then  a  mixture  of  two  volumes  of  hydro- 
gen and  one  of  oxygen  is  introduced.     A  spark  from  an 

induction  coil  is  then  caused 
to  pass  between  the  two 
platinum  wires  in  the  limb 
of  the  tube  bearing  the 
stop-cock,  the  thumb  being 
held  tightly  over  the  open 
end  of  the  tube  at  the  same 
time.  There  is  an  explo- 
sion, and  the  mercury  rises 
and  apparently  fills  the  tube 
completely.  Both  gases  have 
disappeared,  and  the  steam 
produced  condenses  to  liquid 
water,  the  volume  occupied  by 
which  is  very  small  by  comparison  with  that  of  the  gases 
which  combine  to  produce  it.  By  a  simple  modification  of 
this  experiment  the  volume  of  the  steam  produced  can 
also  be  measured.  All  that  is  necessary 
is  to  heat  the  U-shaped  tube  to  a  temper- 
ature above  that  at  which  the  steam  con- 
denses. This  is  accomplished  by  sur- 
rounding it  with  a  wider  tube  and  pass- 
ing, through  the  space  between,  the  vapor 
of  some  liquid  which  boils  at  a  higher 
temperature  than  water.  Amyl  alco- 
hol is  a  suitable  liquid.  The  apparatus  is 
represented  in  Fig.  5.  The  result  shows 
that  the  volume  of  the  steam  is  equal  to 
that  of  the  hydrogen.  We  have  already 
seen  that  this  is  double  the  volume  of 
the  oxygen.  Hence,  2  volumes  hydro-  FIG.  s.— synthesis  of 
gen  -f-  1  volume  oxygen  =  2  volumes 

Steam.  steam. 


WATER  9 

Thus  the  relation  by  volume  of 

Hydrogen :  steam     =  2:2  or  1:1. 
Oxygen :  steam        =1:2. 
Hydrogen :  oxygen  =  2:1. 

It  is  important  to  notice  how  very  simple  these  rela- 
tions are. 

11.  Composition  of  water  by  weight. — When  once  we 
know  the  composition  of  water  by  volume  it  is  easy  to  calcu- 
late its  composition  by  weight.  The  specific  gravity  of  oxy- 
gen referred  to  hydrogen  is  nearly  16.  This  means  that  a 
quart  or  a  liter  of  oxygen  will  weigh  16  times  as  much  as  a 
quart  or  a  liter  of  hydrogen.  Let  us  suppose  that  two  liters 
of  hydrogen  have  combined  with  one  liter  of  oxygen  produ- 
cing, of  course,  two  liters  of  steam.  If  we  call  the  weight  of 
one  liter  of  hydrogen  1,  then  2  parts  of  hydrogen  by  weight 
have  entered  the  compound.  If  oxygen  is  16  times  as  heavy, 
the  liter  of  oxygen  will  weigh  16.  Thus,  2  parts  of  hydro- 
gen +  16  parts  of  oxygen  disappear,  and,  since  no  weight 
is  ever  lost  or  gained  in  chemical  changes,  18  parts  of  water 
must  be  produced.  Water,  therefore,  contains  -fg  or  ^  of  its 
weight  of  hydrogen  and  f  of  oxygen. 

This  calculation  from  the  volumes  of  the  two  gases 
is  sufficient,  but  we  may  corroborate  the  result  by  a  method 
which  proceeds  purely  by  weight.  Several  compounds  of 
copper  and  oxygen  exist,  of  which  the  most  important  is 
black,  and  is  called  cupric  oxide.  When  it  is  heated  in  an 
atmosphere  of  hydrogen  it  behaves  like  many  other  oxides; 
the  hydrogen  removes  the  oxygen,  producing  water  with  it, 
and  copper  as  metal  is  left.  An  oxide  is  always  a  compound 
of  oxygen  with  some  other  element,  and  we  say  that  the 
oxide  is  reduced  by  the  hydrogen.  Now,  if  we  have  weighed 
beforehand  the  vessel  containing  the  cupric  oxide,  and  if  we 
weigh  it  again  after  the  experiment,  it  is  clear  that  the 
loss  in  weight  is  oxygen,  which  has  been  removed  and  has 


10  ELEMENTARY  CHEMISTRY 

produced  water  with  the  hydrogen.  If  we  have  arranged 
matters  so  that  the  water  produced  has  all  been  collected 
in  another  weighed  vessel,  or  in  a  number  of  vessels,  all  of 
which  have  been  weighed  together,  the  increase  in  weight 
of  this  will  be  the  amount  of  water  produced.  And,  since 
the  water  contains  only  two  things,  the  weight  of  the  hydro- 
gen can  be  obtained  by  subtracting  that  of  the  oxygen  from 
that  of  the  water.  The  result  of  this  process  confirms  that 
of  the  other,  and  informs  us  that  water  contains  about  8 
parts  of  oxygen  to  1  of  hydrogen  by  weight. 

This  problem  of  determining  the  proportions  in  which 
hydrogen  and  oxygen  exist  in  water  is  an  important  one  in 
our  science,  and  recently  much  careful  work  has  been  done 
upon  it.  It  has  been  found  that  the  figure  8  is  too  high,  and 
that  water  contains  7.94  parts  of  oxygen  to  1  of  hydrogen. 
This  figure  is  probably  correct  to  the  second  decimal  place. 

It  is  important  to  notice  that  the  composition  of  water 
is  always  the  same.  It  can  be  obtained  in  nature  in  various 
ways.  We  can  distil  ocean  water,  or  melt  ice,  or  condense 
moisture  from  the  air  by  cooling,  and  we  can  prepare  it 
artificially  from  other  substances  by  hundreds  of  different 
methods,  for  water  appears  more  frequently  as  a  product 
of  chemical  changes  than  any  other  substance.  But  when, 
at  last,  we  obtain  pure  water,  it  is  always  found  to  consist 
of  hydrogen  and  oxygen  in  the  same  proportions. 

12.  Decomposition  of  water  by  heat. — We  have  seen  that 
at  100°  liquid  water  passes  into  steam,  the  pressure  being 
one  atmosphere,  and  that  it  is  impossible  to  cause  the  tem^ 
perature  to  rise  above  100°  until  all  the  liquid  has  disap- 
peared. Of  course,  when  the  liquid  is  all  gone,  it  is  possible 
to  heat  the  steam  to  any  desired  degree.  It  expands  like  all 
gases  and  vapors  until,  at  a  yellow  heat,  it  begins  to  separate 
into  its  elements.  The  decomposition  does  not  become  com- 
plete :  steam,  oxygen,  and  hydrogen  are  found  together  in 
the  vessel.  If  the  temperature  is  raised  further,  more  of  the 


WATER  11 

steam  in  the  vessel  is  decomposed,  and  the  higher  the  tem- 
perature the  less  steam  and  the  more  hydrogen  and  oxygen 
are  found  in  it.  At  the  highest  temperatures  at  which  ex- 
periments have  been  made  about  half  of  the  steam  is  split 
up  into  its  constituents,  and  this  is  more  than  1,000°  above 
the  point  at  which  the  decomposition  begins.  No  doubt 
there  is  some  very  elevated  temperature  at  which  the  decom- 
position becomes  complete.  This  we  know,  because  many 
similar  cases  have  been  investigated  in  which  the  tempera- 
tures are  not  so  high,  and  investigation  not  so  difficult. 

Finally,  why  do  we  consider  ice,  liquid  water,  and  water- 
vapor  as  three  forms  of  the  same  substance  ?  Chiefly  because 
we  can  convert  any  one  of  the  three  into  any  other  without 
adding  or  subtracting  anything  except  energy.  Subtract 
heat-energy  from  water  and  it  becomes  ice ;  add  energy  to  it 
and  it  becomes  steam.  In  winter  the  water-vapor  of  the  air 
yields  up  its  heat-energy  to  cold  window-panes  and  becomes 
ice  directly,  without  passing  through  the  liquid  stage. 


CHAPTER  II 

SOLUTION 

13.  Solubility  and  insolubility. — A  little  coarsely  pow- 
dered rosin  is  placed  in  a  beaker,  covered  with  water,  and 
stirred.    There  is  no  apparent  effect.     The  water  is  poured 
off,  and  alcohol  added  instead.   On  stirring,  the  rosin  disap- 
pears, and  it  is  found  that  every  drop  of  the  alcohol  contains 
it.    It  will  be  seen  that  rosin  behaves  differently  with  water 
and  with  alcohol,  and  we  express  this  by  saying  that  it  is 
insoluble  in  the  water  and  soluble  in  the  alcohol.    A  liquid 
which  has  acted  upon  a  substance  soluble  in  it  and  has  taken 
it  up  in  this  way  is  called  a  solution',  thus  we  now  have  an 
alcoholic  solution  of  rosin. 

Alcohol  dissolves  many  solids  which  are  insoluble  in 
water — for  instance,  many  substances  of  animal  and  vege- 
table origin;  but  water  also  dissolves  many  substances  in- 
soluble in  alcohol — for  instance,  numerous  compounds  of 
the  metals,  metallic  salts,  as  they  are  termed.1  Thus, 
when  cupric  sulphate — a  compound  of  copper,  sulphur,  and 
oxygen — is  stirred  up  with  alcohol  it  is  unaffected,  but 
when  treated  with  water  it  dissolves,  communicating  to  the 
liquid  a  blue  color. 

14.  Saturated  solutions, — It  is  worth  noting  that  any 
cupric  salt  which  dissolves  in  water  will  communicate  to  it 

1  Tt  must  be  pointed  out  that  the  term  insoluble  represents  a  state 
of  things  which  probably  is  never  really  met  with.     There  is  good  rea- 
son to  think,  for  instance,  that  there  is  no  substance  insoluble  in  water; 
but  the  solubility  of  some  is  so  very  small  that  it  can  not  be  detected. 
12 


SOLUTION  13 

the  same  color.  In  the  same  way  all  cobalt  salts  are  red  in 
solution,  all  nickel  salts  green,  and  so  on.  Now  let  more 
cupric  sulphate  be  added,  little  by  little,  to  the  solution,  stir- 
ring each  time  until  the  solid  has  disappeared  before  adding 
the  next  portion.  Finally,  there  comes  a  time  when  the 
liquid  refuses  to  dissolve  any  more  cupric  sulphate,  and  the 
latter  remains  unaltered  at  the  bottom.  The  solution  is 
now  saturated.  Thus  a  saturated  solution  is  a  liquid  which 
has  dissolved  all  of  a  given  solid  that  it  will  dissolve  at 
some  temperature  which  is,  or  ought  to  be,  stated.  One 
can  only  speak  of  a  solution  as  being  saturated  with  re- 
spect to  some  particular  solid;  otherwise  the  term  has  no 
meaning.  Our  solution,  which  is  saturated  with  cupric 
sulphate,  is  still  able  to  dissolve  other  substances.  And 
the  temperature  must  be  stated,  because  the  solubility  of 
all  substances  in  all  liquids  is  different  at  different  tem- 
peratures. 

15.  Effect  of  temperature  on  the  amount  dissolved. — Let 
us  heat  the  liquid  which  still  contains  some  undissolved 
cupric  sulphate.  The  solid  dissolves,  and  if  the  liquid  is 
heated  just  short  of  boiling,  it  will  dissolve  about  five  times 
as  much  cupric  sulphate  as  at  the  ordinary  temperature  of 
the  room. 

If  the  solution  is  now  allowed  to  cool,  about  four-fifths 
of  the  cupric  sulphate — that  is,  the  excess  above  what  the 
water  can  dissolve  at  the  ordinary  temperature — will  sepa- 
rate in  crystals. 

This  is  the  usual  behavior  of  solids;  their  solubility  in 
liquids  is  almost  always  increased  by  heating.  There  are  a 
few  exceptions.  The  well-known  substance  gypsum,  or 
plaster  of  Paris,  is  more  soluble  in  water  at  the  ordinary 
temperature  than  in  boiling  water.  And  there  are  sub- 
stances whose  solubility  is  not  much  affected  by  change  of 
temperature.  Thus  table-salt  is  only  slightly  more  soluble 
in  boiling  water  than  in  water  at  the  freezing-point. 


14  ELEMENTARY  CHEMISTRY 

16.  Supersaturated    solutions. — Sodium    sulphate    is    a 
white  compound  of  sulphur  and  oxygen  with  the  metal  so- 
dium.   Some  water  is  warmed  to  about  blood  heat  (37°)  and 
sodium   sulphate   gradually   added.      The   water   dissolves 
about  one-half  its  weight  of  the  solid.    The  liquid  is  freed 
from  undissolved  material  by  pouring  it  through  a  disk  of 
porous  paper  placed  in  a  funnel,  and  is  allowed  to  run  into 
a  clean  bottle,  which  is  immediately  corked,  to  prevent  the 
access  of  dust.     It  will  now  cool  to  room  temperature,  say 
20°,  without  the  separation  of  any  of  the  dissolved  substance. 
But  the  saturated  water  solution  of  sodium  sulphate  at  20° 
only  contains  one-fifth  of  its  weight  of  the  solid.    Our  solu- 
tion, which  contains  nearly  three  times  as  much,  is  said  to 
be  supersaturated. 

A  supersaturated  solution,  therefore,  dontains  more  of 
a  substance  dissolved  than  the  saturated  solution  at  the 
same  temperature.  It  can  only  exist  in  the  absence  of  the 
solid  which  tends  to  separate  from  it.  A  crystal  of  sodium 
sulphate  is  thrown  into  the  bottle  containing  the  supersatu- 
rated solution.  At  once  the  liquid  is  filled  with  crystals  of 
sodium  sulphate  which  are  wet  with  the  saturated  solution 
of  the  salt. 

Most  substances  yield  supersaturated  solutions  to  some 
extent,  but  some  do  so  much  more  easily  than  others,  and 
it  is  found  that  those  substances  which  produce  large,  well- 
formed  crystals  yield  supersaturated  solutions  more  easily 
than  those  which  crystallize  imperfectly. 

SOLUTIONS  OP  LIQUIDS  IN  LIQUIDS 

17.  Some  water  is  placed  in  a  stoppered  funnel  with  a 
stop-cock  at  the  bottom — a  so-called  separating  funnel.    To 
this  ether  is  addefl,  little  by  little,  the  funnel  being  shaken 
after  each  addition.    The  ether  disappears,  entering  into  so- 
lution in  the  water.    Upon  continuing  the  addition  of  ether 
the  water  becomes  saturated  with  it,  and  there  separates, 


SOLUTION  15 

above  the  solution  of  ether  in  water,  a  second  layer.  It  is 
usual  to  call  this  layer  "  ether/7  but  this  is  a  loose  state- 
ment. It  is  a  saturated  solution  of  water  in  ether.  The 
same  two  layers  might  have  been  produced  by  putting  the 
ether  in  the  funnel  and  adding  water  gradually  to  it,  but 
in  this  case  the  second  layer  would  have  formed  below,  since 
the  solution  of  ether  in  water  is  heavier  than  the  other. 

A  liquid,  then,  can  dissolve  another  liquid,  just  as  it  can 
dissolve  a  solid,  and  just  as  in  the  case  of  the  solid  the 
amount  dissolved  in  the  saturated  solution  at  a  fixed  tem- 
perature is  always  the  same.1  But  there  is  a  case  with 
liquids  which  we  do  not  find  with  solids.  Some  liquids  ap- 
pear to  be  able  to  dissolve  each  other  to  any  extent.  Thus 
alcohol  and  water  can  be  mixed  in  any  quantity  whatever, 
and  under  no  circumstances  will  two  layers  be  produced.  It 
has  been  said  that  with  solids,  heating  almost  always  in- 
creases the  amount  which  a  given  liquid  will  dissolve. 
With  liquids  the  exceptions  to  this  are  much  more  nu- 
merous. 

SOLUTIONS  OF  GASES  IN  LIQUIDS 

18.  Some  water  is  placed  in  a  thin-walled  glass — called  a 
beaker — and  heated.  Long  before  the  water  boils  bubbles 
of  gas  collect  on  the  walls  and  finally  rise  to  the  surface  and 
escape.  The  water  has  dissolved  gases  from  the  air.  These 
gases  are  driven  out  by  boiling,  and  then  the  water  loses  its 
pleasant  flavor  and  acquires  the  stale  insipidity  which  is 
familiar  in  boiled  water.  It  will  be  seen  from  this  that  the 
solubility  of  the  gases  of  the  air  in  water  is  reduced  by 
heating  the  water.  This  is  true  of  the  solubility  of  all 
gases  in  all  liquids.  The  solubility  of  the  atmospheric 
gases  in  water  is  slight.  Thus  of  nitrogen,  which  makes 

1  In  both  cases  the  amount  dissolved  also  varies  slightly  with  the 
pressure  on  the  surface  of  the  liquid.  Under  ordinary  circumstances 
the  pressure,  that  of  the  atmosphere,  can  be  considered  constant. 


16  ELEMENTARY  CHEMISTRY 

up  about  four-fifths  of  the  air  by  volume,  water  dissolves 
less  than  one-fiftieth  of  its  bulk  at  ordinary  temperature  and 
pressure.  Some  gases  dissolve  very  freely.  Water  at  the 
freezing-point  dissolves,  for  instance,  over  eleven  hundred 
times  its  volume  of  ammonia  gas,  and  this  solution  consti- 
tutes the  water  of  ammonia  of  commerce.  The  law  that 
solubility  of  gases  decreases  with  rising  temperature  is  well 
illustrated  in  this  case.  When  the  solution  is  heated  to 
boiling  all  but  the  merest  trace  of  ammonia  escapes. 

19.  Effect  of  pressure  on  the  solubility  of  gases. — It  has 
been  seen  that,  in  the  case  of  solids,  the  quantity  dissolved 
is  affected  by  pressure,  but  that  the  effect  is  slight  and 
can  be  neglected.    With  gases,  on  the  other  hand,  pressure 
is  an  important  element.     When  the  solubility  of  a  gas 
in  a  liquid  is  slight  the  quantity  dissolved  is  proportional 
to  the  pressure. 

By  pressure  is  meant  not  the  total  pressure  on  the 
surface  of  the  liquid,  but  that  fraction  of  it  which  is  exerted 
by  the  gas  under  discussion. 

Carbon  dioxide,  a  compound  of  carbon  and  oxygen,  is 
a  gas  of  which  water,  at  ordinary  temperature  and  the 
pressure  of  one  atmosphere,  dissolves  about  its  own  vol- 
ume. When  the  pressure  is  doubled  the  quantity  dissolved 
is  doubled.  Soda-water  is  water  into  which  about  three  vol- 
umes of  carbon  dioxide  have  been  pumped  under  a  pressure, 
therefore,  of  about  three  atmospheres.  When  the  liquid 
is  exposed  to  the  air  the  gas  escapes  slowly  with  efferves- 
cence. The  same  state  of  things  exists  in  effervescent  fer- 
mented liquors,  only  here  the  gas,  instead  of  being  pumped 
into  the  liquid,  is  produced  in  the  liquid  by  fermentation  in 
corked  bottles,  and,  since  it  has  no  escape,  dissolves  in  the 
liquid  under  its  own  pressure. 

20.  Supersaturated  solutions  of  gases. — Supersaturated 
solutions  are  common  with  gases.     An  ordinary  glass  of 
plain  soda-water  is  an  excellent  example.    As  it  stands  on 


SOLUTION  17 

the  table  the  air  exerts  on  the  upper  surface  a  pressure 
of  1  atmosphere.  But  the  quantity  of  carbon  dioxide  dis- 
solved does  not  depend  upon  this  total  pressure,  but 
only  upon  that  fraction  of  it  which  is  due  to  carbon 
dioxide. 

Now,  since  the  quantity  of  carbon  dioxide  in  the  air  is 
very  small,  the  pressure  which  this  gas  exerts  on  the  liquid 
is  practically  zero.  Hence  all  the  gas  should  escape.  But 
as  a  matter  of  fact  the  escape  is  extremely  slow,  and  if  the 
glass  is  clean  and  the  liquid  undisturbed,  gas  is  only  given 
off  from  the  surface. 

It  has  been  found  possible  to  pump  carbon  dioxide  into 
a  glass  of  water  under  a  pressure  of  three  atmospheres,  and 
then  remove  the  pressure  without  any 
bubbling.  Of  course  in  this  case  the 
glass  was  carefully  cleaned  and  the 
liquid  absolutely  undisturbed. 

Supersaturated  solutions  of  solids 
crystallize  on  throwing  in  a  fragment 
of  the  substance  which  tends  to  sep- 
arate. Other  solids  have  no  effect. 
This  is  an  interesting  difference  be- 
tween Solids  and  gases.  A  SUpersatu-  PlG-  6. -Supersaturated  aolu- 
,-,,,.  _  ill!  tion  of  a  gas. 

rated  solution  of  a  gas  forms  bubbles 

— gives  off  the  excess — in  contact  with  any  gas  whatever, 
preferably  with  some  other  gas.  Thus,  let  a  perfectly 
clean  glass  tube,  sealed  at  one  end,  be  lowered,  open  end 
own,  into  one  glass  of  soda-water.  Bubbles  ascend  from 
e  plane  of  contact  between  air  and  liquid  (Fig.  6). 

When  a  piece  of  charcoal  is  dropped  into  the  glass  ener- 
getic effervescence  occurs.  The  action  is  the  same  as  in 
the  glass  tube.  Charcoal  is  porous  and  full  of  air,  and  it 
acts  as  a  means  by  which  air  can  be  brought  into  contact 
with  the  liquid. 

It  is  easy  to  remove  the  air  by  heating  the  charcoal  red- 


18  ELEMENTARY  CHEMISTRY 

hot  and  holding  it  under  mercury  to  cool,  and  then,  when 
dropped  into  the  solution  of  carbon  dioxide,  no  bubbles  are 
produced. 

SOLUTIONS  IN  SOLIDS 

21.  Even  when  the  valve  of  a  bicycle  tire  is  tight  and  the 
tire  free  from  punctures,  the  air  escapes  slowly  and  the  tire 
becomes  soft.  It  is  often  said  that  the  tire  is  porous;  but 
this  is  only  another  way  of  saying  that  in  some  manner  the 
air  passes  through  it,  and  does  not  really  help  us  to  under- 
stand the  matter.  Probably  most  people  think  of  the  tire  as 
containing  openings  too  small  to  be  seen,  through  which  the 
air  gradually  streams. 

The  air  consists,  for  the  most  part,  of  two  gases,  one 
of  which,  oxygen,  has  already  been  mentioned.  The  other 
is  also  colorless  and  odorless,  and  is  called  nitrogen.  Now, 
when  a  tire  has  become  soft,  by  standing,  it  is  found  that 
the  gas  which  it  still  contains  is  nearly  pure  nitrogen.  It 
is  the  oxygen  which  has  escaped  most  easily.  This  is  not 
to  be  explained  by  the  assumption  that  oxygen  passes 
through  small  orifices  more  easily  than  nitrogen,  for,  from 
other  experiments,  we  know  that  the  reverse  is  true. 

How,  then,  is  the  behavior  of  the  rubber  tire  to  be  ex- 
plained? By  far  the  simplest  statement  to  make  is  that 
the  rubber  tire  has  dissolved  the  oxygen.  We  have  seen 
that  the  solubility  of  a  gas  is  proportional  to  the  pressure 
of  that  particular  gas  against  the  dissolving  substance.  In- 
side the  tire  the  pressure  of  the  oxygen  is  very  much  greater 
than  outside.  Hence  the  oxygen  continually  dissolves  at 
the  inner  surface  and  is  given  off  at  the  outside. 

In  the  same  way  rubber  will  dissolve  carbon  dioxide. 
Many  facts  similar  to  the  behavior  of  the  pneumatic  tire 
are  known,  and  they  can  all  be  explained  in  the  same  simple 
way.  Thus,  when  hydrogen  is  passed  into  a  hot  platinum 
tube,  it  escapes  rapidly  through  the  walls  of  the  tube.  It  will 
also  pass  through  a  thin  plate  of  iron,  even  in  the  cold. 


SOLUTION  19 

22.  Solids  can  also  dissolve  liquids. — Every  apparently 
dry  powder  which  has  been  exposed  to  the  air  has  obtained 
from  it  some  water.     The  presence  of  the  water  can  be 
proved  by  heating  the  powder  in  a  glass  tube,  when  the  water 
condenses  in  drops  in  the  cool  portion,  and  the  quantity  can 
be  ascertained  by  weighing  a  portion  of  the  powder,  heating 
it  gently,  and  then  reweighing.    The  difference  is  the  weight 
of  the  expelled  water.     In  this  way  we  discover  that  the 
quantity  of  the  water — called  hygroscopic  water — depends 
upon  the  pressure  of  the  water-vapor  in  the  air,  and  is 
greater  on  wet  days  than  on  dry  ones. 

When  a  cylinder  of  lead  is  placed  with  its  lower  sur- 
face on  a  disk  of  gold,  gold  dissolves  in  the  solid  lead  and 
travels  upward  in  the  cylinder,  very  slowly  at  ordinary 
temperatures,  and  more  rapidly  when  heated.  In  the  same 
way  gold  will  travel  into  silver  and  platinum. 

From  these  last  cases  it  will  be  seen  that  solids  can  dis- 
solve each  other.  Glass  is  an  excellent  example  of  a  solu- 
tion, formed  at  a  red  heat  in  the  melting-pot,  which  solidifies 
without  any  separation,  and  is  therefore  to  be  called  a 
solid  solution.  Very  old  glass  is  sometimes  dim  and  almost 
opaque  because  the  constituents  have  separated  in  crystals, 
just  as  copper  sulphate  crystallizes  from  water  when  the  hot 
saturated  solution  is  cooled. 

SOLUTIONS  IN  GASES  AND  VAPORS 

23.  Bromine  is  a  heavy  reddish-black  liquid  which  gives 
off,  at  ordinary  temperatures,  a  suffocating  red  vapor.  When 
a  small  quantity  of  it  is  placed  in  a  strong  glass  tube  it 
evaporates  until  every  cubic  centimeter  of  the  space  in  con- 
tact with  the  liquid  contains  a  certain  quantity  of  the  vapor. 
The  color  of  the  vapor  will  serve  as  a  rough  measure  of  this 
quantity.    Now  let  air  be  forced  into  the  tube  under  strong 
pressure.     At  once  the  color  of  the  vapor  becomes  more 
intense,  and  if  we  have  taken  a  small  enough  quantity  of 

3 


20  ELEMENTARY  CHEMISTRY 

bromine  the  liquid  in  the  bottom  disappears.  By  forcing 
in  enough  air  the  color  may  be  made  six  times  as  strong 
as  before,  which  means  that  six  times  as  much  bromine  has 
evaporated.  There  is  no  reason  why  this  state  of  things 
should  not  be  described  by  saying  that  the  compressed  air 
has  dissolved  the  bromine.  Iodine  and  camphor  behave  in 
a  similar  way  when  compressed  with  different  gases.  Va- 
pors seem  to  have  the  same  power.  Potassium  iodide  is  a 
white  solid  compound  of  a  metal,  potassium,  with  iodine. 
It  will  not  evaporate  below  a  red  heat.  It  is  soluble  in 
alcohol,  and,  when  the  solution  is  enclosed  in  a  strong 
glass  tube  and  heated,  the  alcohol  becomes  a  vapor,  but  the 
solid  still  remains  dissolved,  for  the  contents  of  the  tube 
remain  perfectly  transparent,  no  crystals  separating. 

24.  Solution  common  to  solids,  liquids,  and  gases. — The 
most  important  result  of  all  this  is  that  there  is  no  reason  to 
assume  that  liquids  are  the  only  bodies  which  can  act  as 
solvents.     Solids,  liquids,  and  gases  all  possess  the  power, 
and  all  of  the  three  can  act  as  dissolved  substance.    But  the 
physical  nature  of  a  liquid  is  such  that  solutions  in  it  are 
far  easier  to  make  and  to  handle,  and  hence  liquid  solutions 
are  more  important. 

SUSPENSION 

25.  Eosin  is  insoluble  in  water  but  soluble  in  alcohol. 
The  alcoholic  solution  is  allowed  to  fall  into  a  large  quantity 
of  water  in  a  cylinder.    The  liquid  is  filled  with  a  white  tur- 
bidity of  finely  divided  solid  rosin.    The  rosin,  in  this  con- 
dition, settles  very  slowly,  and  is  said  to  be  suspended  in 
the  water,  and  the  state  of  things  before  it  has  settled 
is  called  a  suspension.     After  a  rain  the  rivers  are  turbid 
with  suspended   matter  which  is  washed  into  them,  and 
which  finally  settles,  leaving  the  water  clear.     If  the  sus- 
pended matter  is  lighter  than  the  rest  of  the  liquid  it  will 
go  to  the  top  instead  of  the  bottom.     Milk  consists  of 


SOLUTION  21 

small  fat  globules  suspended  in  an  aqueous  liquid,  and,  when 
it  is  allowed  to  stand,  they  collect  at  the  surface,  producing 
the  cream. 

26.  Suspensions  in  gases. — Gases  hold  suspended  matter 
in  exactly  the  same  way.    The  air  is  the  most  familiar  exam- 
ple of  this.    It  always  contains  dust,  finely  divided  iron  from 
meteorites,  and  the  bodies  of  bacteria.     A  mist  consists  of 
fine  suspended  water  globules.     Suspended  salt  crystals  are 
carried  by  the  winds  to  points  hundreds  of  miles  from  the 
coast. 

27.  Distinction   between   solution   and   suspension. — It 
will  be  seen  that  a  suspension  is  a  temporary  condition  pro- 
duced by  outside  action,  as,  for  instance,  by  shaking  up  a 
powdered  solid  with  a  liquid.    But  solution  is  a  process  which 
takes  place  of  itself,  and  the  result  is  a  permanent  condition, 
the  natural  state  of  the  two  substances  when  in  contact. 
Thus,  let  some  cupric  sulphate  be  placed  in  a  tall  glass  cylin- 
der, and  water  carefully  poured  over  it  so  as  not  to  disturb 
it.    Then  let  the  cylinder  be  closed,  to  prevent  evaporation, 
and  stood  away  in  a  place  as  free  from  changes  of  tempera- 
ture as  possible.    Since  the  cupric  sulphate  lies  at  the  bot- 
tom, it  is  clear  that  it  is  heavier  than  the  water.    Neverthe- 
less, it  gradually  diffuses  through  the  water,  and  if  there  is 
enough  water  present  the  final  result  is  a  uniform  mixture 
of  the  two.     This  is  permanent.     The  cupric  sulphate  will 
never  settle  to  the  bottom  as  the  rosin  did  in  the  other  case. 
It  is  usually  said  that  a  suspension  differs  from  a  solution 
in  this  respect,  that  the  solid  substance  can  be  removed  by 
filtering  from  the  first,  but  not  from  the  second.     Now,  it 
is  true  that  ordinary  suspended  matter  can  be  removed  by 
filtering,  while  dissolved  matter  never  can.    But  any  filter, 
whether  of  stone  or  paper,  is  simply  a  substance  full  of 
small  holes  which  allow  the  liquid  to  pass  and  retain  the 
solid.     Plainly,  the  behavior  of  any  suspension  toward  a 
filter  depends  upon  whether  the  holes  in   the  filter  are 


22  ELEMENTARY  CHEMISTRY 

smaller  than  the  suspended  particles,  or  the  reverse.  If 
smaller,  the  matter  will  be  removed  and  the  liquid  will  be 
clear.  If  larger,  it  will  pass  through.  It  is  easy  to  imagine 
and  easy  to  obtain  a  suspension  so  fine  that  it  will  pass  with- 
out change  through  any  filter  which  can  be  made  at  present. 
Thus  gold,  silver,  platinum,  and  other  metals  have  recently 
been  obtained  in  exceedingly  fine  suspension  in  water.  The 
liquids  are  dark  in  color  and  appear  perfectly  clear,  and  the 
solid  particles  are  so  small  that  they  are  not  visible  under 
the  most  powerful  microscope.  But  they  are  suspensions, 
as  is  shown  by  the  fact  that  the  metal  separates,  after  a 
time,  as  well  as  by  optical  methods  which  can  not  be  dis- 
cussed here. 


CHAPTER  III 
PHYSICAL  AND  CHEMICAL  CHANGE 

THE  LAW  OF  THE  INDESTRUCTIBILITY  OF  MATTER 

28.  Changes  in  matter. — The  material  bodies  which  sur- 
round us  are  subject  to  continual  change.     Water  left  in 
free  air  disappears  and  enters  the  atmosphere  as  an  in- 
visible vapor,  from  which  it  appears  again  as  solid — snow 
or  hail,  or  as  liquid — rain,  mist,  or  cloud,  according  to  the 
conditions.     When  cooled  it  solidifies  as  ice  and  acquires 
quite  different  properties.     Iron  left  exposed  to  air  and 
water  passes,  first  at  the  surface,  and  at  last  completely, 
into  a  soft,  yellowish-red  substance  quite  different  from  the 
iron,  which  we  know  under  the  name  of  rust.    Bright  sur- 
faces of  metals  tarnish  with  a  speed  which  depends  upon 
the  nature  of  the  metal  and  of  the  air  with  which  it  is  in 
contact.    Even  the  rocks,  so  often  referred  to  as  examples 
of  permanence,  crumble  under  the  influence  of  the  weather 
into  earth  or  clay,  which  is  finally  carried  down  into  the 
sea  and  made  into  rock  again.     More  rapid  changes,  like 
those  which  take  place  in  a  coal  fire,  or  in  a  burning  candle, 
are  familiar.    Finally,  the  very  complicated  changes  which 
take  place  in  the  living  body  must  be  mentioned. 

29.  Physical  and  chemical  change. — It  is  convenient  to 
divide  changes  of  the  kind  referred  to  into  two  classes,  the 
nature  of  which  will  be  clear  from  a  few  examples. 

PHYSICAL  CHANGES 

30.  Sulphur. — A  roll  of  sulphur  is  brought  near  to  a 
small  ball  of  pith  or  cork  which  hangs  by  a  thread.    There  is 

23 


24  ELEMENTARY  CHEMISTRY 

no  result.  Now  let  the  roll  be  rubbed  with  a  cloth,  and  again 
brought  near  the  ball.  This  time  the  ball  is  attracted  and 
swings  far  out  of  the  natural  position.  By  friction  the 
sulphur  received  an  electric  charge  and  this  conferred  upon 
it  new  properties,  but  it  is  still  sulphur.  In  a  little  while  the 
charge  is  lost  and  the  new  properties  disappear. 

31.  Mercuric  iodide. — Mercuric  iodide  is  a  bright-red, 
heavy  powder  composed  of  small  crystals.  It  is  a  chemical 
compound  of  the  familiar  liquid  metal  mercury — used  in 
filling  thermometers — with  iodine,  a  grayish-black  sub- 
stance looking  like  graphite.  Some  mercuric  iodide  is  placed 
in  a  tube  and  gently  heated.  Slightly  above  the  boiling-point 
of  water  a  striking  change  of  color  to  bright  yellow  occurs, 
and  it  can  be  shown  that  the  shape  of  the  small  crystals,  of 
which  the  powder  consists,  has  changed.  After  having  com- 
pared the  yellow  with  the  red  modification — the  two  have 
exactly  the  same  composition,  both  containing  200  parts  of 
mercury  and  254  parts  of  iodine — the  yellow  powder  is  di- 
vided into  two  portions.  Half  of  it  is  placed  on  paper  and 
rubbed  briskly  with  a  glass  rod.  At  once  it  changes  back 
again  to  the  red  modification.  The  other  half  is  preserved 
in  a  clean  tube  and,  after  some  time,  examined.  The  same 
change  has  occurred:  the  yellow  powder  is  now  red.  Evi- 
dently the  only  result  of  the  pressure  of  the  glass  rod  was 
to  accelerate  a  change  which  would  have  occurred  without 
it  if  sufficient  time  was  given. 

At  ordinary  temperatures  the  natural  state  into  which 
mercuric  iodide  passes  if  left  to  itself  is  the  red  modifica- 
tion; at  slightly  higher  temperatures,  accurately  above  129°, 
the  yellow  modification  is  the  natural  state,  and  the  red 
passes  into  it.  We  express  this  by  saying  that  below  129° 
the  red  modification  is  stable,  above,  the  yellow.  There  is 
much  in  all  this  which  reminds  us  of  the  change  of  ice  into 
water  and  back.  Ice  and  liquid  water,  like  yellow  and  red 
mercuric  iodide,  are  two  modifications  of  the  same  substance. 


PHYSICAL  AND  CHEMICAL  CHANGE  25 

Below  0°  ice  is  the  stable  modification;  above  0°  water.  The 
most  important  difference  is  that,,  in  the  case  of  water,  one 
modification  is  liquid  and  the  other  solid,  while  in  the  case 
of  mercuric  iodide  they  are  both  solid. 

32.  Tin. — The  properties  of  all  material  bodies  depend 
upon  the  circumstances  under  which  they  exist.    The  most 
important  of  these  are  pressure  and  temperature,  and  it  is 
only  because  pressure  and  temperature  do  not  vary  widely 
in  ordinary  life  that  we  are  in  the  habit  of  regarding  the 
properties  of  bodies  as  tolerably  constant.     In  order  to 
fix  our  ideas,  let  us  think  of  a  cube  of  tin.     It  is  clear, 
first  of  all,  that  the  volume  of  the  cube  depends  upon  the 
pressure  and  temperature.    With  rising  temperature  it  will 
expand,  with  increasing  pressure  it  will  contract.    Whether 
it  is  a  solid,  a  liquid,  or  a  vapor  is  chiefly  a  matter  of  tem- 
perature.   Slight  heating  easily  converts  it  into  a  metallic, 
mercury-like  liquid,  and,  on  intense  heating,  this  can  be 
transformed  into  a  vapor.     On  exposing  the  cube  to  win- 
ter's cold  for  a  long  time,  it  falls  to  a  loose,  dull-gray  mass, 
which  is  simply  another  solid  modification  of  tin,  and  is 
reconverted  into  the  familiar  white  lustrous  metal  at  once 
on  heating — for  instance,  by  pouring  hot  water  over  it. 
The  vapor,  the  liquid,  and  the  two  solids  are  all  tin;  no 
new  substance  has  been  produced. 

Changes  like  the  variation  in  the  electrical  state  of  sul- 
phur, the  transformation  of  water  into  ice,  the  change  of  the 
red  mercuric  iodide  into  the  yellow  and  back,  the  production 
of  the  different  modifications  of  tin,  all  agree  in  this  re- 
spect, that  no  new  substance  is  formed  anywhere  in  the 
process.  Taken  together  they  are  called  physical  changes, 
and  their  study  is  the  subject-matter  of  the  science  of 
physics. 

CHEMICAL  CHANGES 

33.  Mercuric   oxide. — This    compound    of   oxygen    and 
mercury  ia  a  dense  red  substance.    A  little  of  it  is  placed  in 


26 


ELEMENTARY  CHEMISTRY 


a  hard  glass  retort/  A,  and  heated.  The  retort  is  connected 
with  a  tube  called  the  delivery  tube.  This  leads  away 
any  gas  which  may  be  evolved  to  a  vessel  where  it  can 
be  collected  in  an  inverted  cylinder  full  of  water,  D  (Fig. 
7).  The  first  effect  of  the  heat  is  to  turn  the  mercuric 
oxide  from  red  to  black.  On  heating  to  a  higher  tempera- 
ture, bubbles  of  a 
colorless  gas  rise 
through  the  water 
and  fill  the  invert- 
ed cylinder,  and  if 
enough  mercuric 
oxide  has  been 
taken,  several  cyl- 
inders full  of  the 
gas  may  be  col- 
lected in  this  way. 
...  .  Upon  investiga- 

Fio.   7.— Decomposition  of   mercuric    oxide    by   heat.  r 

A,  retort    in   which   the   su&stance  is  heated  ;   B,  tion        the        g&S 

receiver  in  which  the  mercury  collects;    D  and  C,  r,rovpq  |o  Kp  Ofjor- 
apparatus  for  collecting  the  oxygen  over  water. 

less  as  well  as  col- 
orless; when  a  match-stick  bearing  a  spark  is  placed  in  a 
tube  full  of  it,  the  wood  bursts  into  flame.  Kemembering  the 
experiment  in  which  we  split  up  water  by  the  electric  cur- 
rent, we  perceive  that  this  gas  is  oxygen. 

Upon  examining  the  retort  it  is  found  that  the  red 
powder  has  disappeared  and  the  retort  contains  a  metallic, 
mirror-like  ring  around  the  upper  cool  portion  and  a  drop 
of  metallic  liquid,  both  of  which  are  easily  seen  to  consist 
of  mercury.  The  heat  has  separated  the  mercuric  oxide 
into  its  constituents,  and  almost  all  compounds  behave  simi- 
larly at  a  sufficiently  high  temperature.  Here,  then,  one 


1  Hard  glass  is  a  variety  which  will  stand  great  heat  without  melt- 
ing and  is  used  when  the  temperature  required  is  high. 


PHYSICAL  AND  CHEMICAL  CHANGE  27 

substance,  mercuric  oxide,  has  disappeared,  and  two  new  sub- 
stances, oxygen  and  mercury,  have  arisen  in  its  place. 
Neither  can  be  regarded  in  any  sense  as  a  form  or  modifi- 
cation of  mercuric  oxide.  This  is  a  change  of  a  different 
character  from  the  change  of  ice  to  water  or  of  red  mer- 
curic iodide  to  yellow. 

34.  Magnesium. — As  another  example,  let  us  burn  a 
piece  of  magnesium  ribbon.    There  is  an  intense  evolution 
of  light  and  heat,  and  the  result  is  a  white  powder  of  mag- 
nesia— magnesium  oxide — which  would  weigh,  if  all  collect- 
ed, considerably  more  than  the  magnesium  burned.  The  oxy- 
gen of  the  air  has  combined  with  the  magnesium  and  has 
produced  the  new  substance,  magnesium  oxide.    The  change 
is  chemical.    A  chemical  change,  then,  is  a  change  in  which 
new  substances  appear.     The  study  of  chemical  changes  is 
the  province  of  chemistry.     For  a  long  time  it  was  held 
that  the  task  of  the  chemist  was  finished  when  the  products 
of  the  change  had  been  investigated  completely.    This  view 
is  now  given  up.     The  investigation  of  the  products  of 
chemical  changes  is  an  important  part  of  chemical  science, 
and  it  is  the  part  with  which  this  book  is  chiefly  concerned, 
but  it  is  only  a  part;  after  this  there  follows  the  study  of 
the  change  itself,  its  speed,  and  the  variations  in  this  speed 
when  temperature,  pressure,  and  other  conditions  alter,  and 
only  when  this  has  been  accomplished  can  we  claim  really 
to  understand  the  matter.1 

THE  LAW  OF  THE  INDESTRUCTIBILITY  OF  MATTER 

35.  Imagine  two  glass  globes  which  have  been  sealed 
tightly,  so  that  nothing  can  leave  and  nothing  enter  them. 

1  It  will  be  an  excellent  exercise  for  the  student  to  classify  changes 
in  material  things  which  occur  in  daily  life  into  chemical  and  phys- 
ical. Thus  it  is  easily  seen  that  the  various  transformations  of  water 
and  the  melting  of  metals  are  physical  phenomena,  while  the  rusting 
of  iron,  the  decay  of  wood,  and  the  burning  of  a  candle  are  chemical. 


28  ELEMENTARY  CHEMISTRY 

Imagine  also  that  the  globes  contain  some  material  in  which 
a  chemical  or  physical  change  can  occur.  As  an  example 
of  the  first  we  may  consider  that  the  first  globe  is  filled 
with  air  and  contains  a  piece  of  magnesium  wire.  As  an 
example  of  the  second  let  us  consider  that  the  second  globe 
contains  some  red  mercuric  iodide.  The  globe  containing 
the  magnesium  wire  is  placed  on  one  side  of  a  delicate  scale 
and  accurately  balanced  by  weights  placed  on  the  other  side. 
The  globe  containing  the  mercuric  iodide  is  treated  in  the 
same  way. 

Now  let  us  cause  the  two  changes  to  occur.  It  will  be 
easy  to  focus  the  rays  of  the  sun  with  a  lens  on  the  mag- 
nesium until  it  takes  fire  and  burns.  The  red  mercuric 
iodide  can  be  warmed  by  the  careful  application  of  a  flame 
to  the  outside  of  the  glass,  until  it  changes  to  the  yellow 
modification.  The  globes  must  now  be  allowed  to  cool,  for 
it  is  impossible  to  weigh  a  hot  object  accurately,  owing  to 
air  currents.  Then  they  are  replaced,  each  in  its  own  scale. 
It  will  be  found  that  the  weight  of  both  is  unaltered.  We 
might  try  any  number  of  changes  in  this  way,  and  the  re- 
sult would  always  be  the  same.  It  is,  in  fact,  impossible 
to  change  the  weight  of  the  globes  except  by  piercing  the 
glass  and  taking  something  out  or  putting  something  in 
through  the  orifice. 

Now  the  weight  measures  the  quantity  of  matter  in 
the  globes  just  as  the  shopkeeper's  weights  measure  the 
quantities  of  various  substances  he  hands  to  his  customers. 
Hence  chemical  and  physical  changes  produce  no  alteration 
in  the  quantity  of  matter  present.  This  fact,  that  nothing 
is  ever  gained  or  lost  in  any  processes  which  man  has  in- 
vestigated, is  usually  called  the  law  of  the  indestructibility 
of  matter.  It  is  often  stated  in  this  form :  "  The  quantity 
of  matter  in  the  universe  is  constant."  But  the  universe  is 
a  large  place,  and  it  is  dangerous  to  make  statements  about 
what  may  happen  in  inaccessible  portions  of  it.  There  is, 


PHYSICAL  AND  CHEMICAL  CHANGE  29 

however,  no  reason  why  we  should  not  extend  the  law  to  our 
own  solar  system.  On  the  sun,  physical  processes  and — in 
the  cooler  portions  of  its  atmosphere — chemical  processes 
also,  take  place  on  a  gigantic  scale.  But  any  considerable 
alteration  in  the  mass  of  the  sun  would  alter  the  pull  which 
it  exerts  on  the  planets,  and  produce  a  modification  in 
their  motions  which  the  modern  telescope  could  not  fail  to 
detect.  That  no  such  change  is  detected  shows  that  the 
mass  of  the  sun  remains  the  same.  Of  course  the  same 
reasoning  applies  to  the  planets  and  their  satellites. 

The  fact  that  the  length  of  the  year  remains  the  same 
shows  that  matter  does  not  appear  out  of  nothing  nor  vanish 
into  nothing  on  either  the  sun  or  the  earth.  The  fact  that 
the  time  of  an  eclipse  can  be  accurately  predicted  centuries 
in  advance  is  a  very  sharp  proof  that  the  weights  of  the  sun, 
the  moon,  and  the  earth  remain  the  same. 

36.  Importance  and  nature  of  the  law. — This  law  is  the 
foundation  of  chemical  and  physical  science.    Like  all  the 
other  laws  of  science,  it  is  simply  a  descriptive  statement 
which  sums  up,  in  a  brief  way,  the  result  of  man's  experi- 
ence.    It  does  not  depend  at  all  upon  any  ideas  we  may 
entertain  regarding  the  structure  of  material  things,  and 
is  equally  true  whether  we  consider  them  as  made  up  of 
particles  with  spaces  between  them  or  not.     On  the  con- 
trary, the  ideas  which  we  hold  upon  such  subjects  must 
be  made  to  agree  with  this  and  other  facts,  to  be  pres- 
ently discussed.     Otherwise  they  are  worthless  and  mis- 
leading. 

37.  Apparent  exceptions. — Cases  which  apparently  con- 
tradict this  law  are  common.    The  disappearance  of  water 
exposed  to  the  air  and  the  burning  away  of  a  candle  with- 
out residue  are  examples.     The  water  is  diffused  through 
the  air  as  an  invisible  vapor.    As  for  the  candle,  it  consists 
essentially  of  hydrogen  and  carbon.    When  it  burns,  both 
combine  with  the  oxygen  of  the  air,  the  hydrogen  to  water- 


30  ELEMENTARY  CHEMISTRY 

vapor,  and  the  carbon  to  the  colorless  gas,  carbon  dioxide, 
already  referred  to.  And,  of  course,  the  sum  of  the  weights 
of  the  candle  burned  and  the  oxygen  taken  from  the  air 
must  be  equal  to  the  sum  of  the  weights  of  the  products  of 
the  burning. 


JOSEPH  PEIESTLEY 
B.  England ,  1733.     D.  Pennsylvania,  1804. 


CHAPTER  IV 
MIXTURE-ELEMENT- COMPO  UND 

THE  LAW  OF  DEFINITE  PROPORTIONS 

38.  Mixtures  and  compounds. — We  have  some  roll  sul- 
phur which  has  been  finely  powdered  and  some  copper  fil- 
ings.   A  small  quantity  of  the  copper  is  mixed  carefully  with 
about  half  its  weight  of  the  sulphur,  and  the  mass  examined. 
It  is  easily  seen  to  consist  of  copper  and  sulphur,  lying  side 
by  side,  and  this  would  at  once  inform  us  that  it  is  a  mix- 
ture of  two  substances.    Now  suppose  the  particles  of  cop- 
per and  sulphur  to  be  so  small  that,  to  the  eye,  the  powder 
appears   uniform.     In  this   case   the   application   of   the 
microscope  would  reveal  the  presence  of  both.    But  suppose 
the  particles  of  both  to  be  so  small  that  no  microscope 
makes  them  visible.    What  then  ?    How  can  we  be  sure  that 
we  have  here  a  mixture  of  two  substances  and  not  a  com- 
pound, which,  of  course,  is  one  substance?     Here  we  are 
driven  to  use  methods  of  separating  the  copper  from  the 
sulphur. 

39.  Separation  of  mixtures. — For  instance,  it  is  easy  to 
obtain  a  liquid  heavier  than  the  sulphur  but  lighter  than  the 
copper.    When  the  mixture  is  thrown  into  this  the  sulphur 
floats  and  the  copper  sinks.    Or  we  can  shake  up  the  mix- 
ture with  water.     Both  sink,  but  the  sulphur,  being  much 
lighter,  remains  suspended  longer,  and  by  pouring  off  at  the 
right  time,  and  repeating  the  "  elutriation,"  as  it  is  called, 
the  two  may  be  separated.     Or,  finally,  we  can  treat  the 
mixture  with  a  liquid  which  dissolves  one  constituent  and 

31 


32  ELEMENTARY  CHEMISTRY 

not  the  other.  Thus  carbon  disulphide,  a  colorless  inflam- 
mable liquid,  dissolves  sulphur  freely  and  copper  not  at  all, 
and  when  the  mixture  is  stirred  with  it  the  sulphur  dissolves 
and  can  be  recovered  by  pouring  off  the  liquid  from  the  cop- 
per and  allowing  it  to  evaporate. 

Now  let  the  mixture  of  copper  and  sulphur  be  placed  in  a 
test-tube,  and  strongly  heated  at  one  point.  There  is  a  con- 
siderable evolution  of  light  and  heat.  When  the  tube  cools 
we  break  it,  and  find  a  bluish-black,  uniform  substance 
called  copper  sulphide,  in  which  no  copper  and  no  sulphur 
can  be  perceived  in  any  way,  and,  applied  to  which,  the 
methods  of  separation,  which  succeed  with  the  mixture, 
fail.  For  instance,  if  we  powder  the  copper  sulphide  and 
throw  it  into  a  liquid  whose  density  lies  between  that  of 
copper  and  sulphur,  no  separation  results.  Copper  sulphide 
is  one  substance,  having  its  own  density,  and  what  happens 
depends  entirely  on  whether  the  copper  sulphide  is  lighter 
or  heavier  than  the  liquid.  If  heavier  it  will  all  sink;  if 
lighter  it  will  all  float.  N"or  will  carbon  disulphide  dissolve 
any  sulphur  from  copper  sulphide. 

It  will  be  seen  that  the  distinction  between  a  mechanical 
mixture  of  two  different  powders  and  a  chemical  compound 
of  the  same  two  substances  is  an  easy  one  to  make.  The 
methods  of  separation  depend  upon  the  fact  that  the  mix- 
ture consists  of  two  solids  lying  side  by  side,  and  that  the 
two  solids  have  different  properties.  The  next  question  is 
whether  there  are  any  bodies  regarded  as  mixtures,  upon 
which  all  these  methods  of  separation  fail ;  in  other  words, 
are  there  mixtures  which  can  not  be  thought  of  as  con- 
sisting of  two  bodies  very  finely  divided  and  mixed — mix- 
tures which  are  not  only  uniform  to  the  eye,  but  really  so. 
The  student  will  at  once  think  of  mixtures  of  gases  like 
the  air,  and  of  the  various  cases  of  solutions  in  liquids  and 
solids  which  have  been  discussed.  These  are  mixtures  which 
are  perfectly  uniform,  and  the  separation  of  which,  by 


MIXTURE-ELEMENT—COMPOUND  33 

methods  like  those  which  succeed  with  copper  and  sulphur, 
is  usually  impossible.  Why,  then,  are  bodies  like  glass, 
air,  and  salt  water  regarded  as  mixtures  at  all?  Why  not 
regard  each  as  one  substance — as  a  compound  ?  The  answer 
to  this  is  simple  and  of  supreme  importance. 

40.  The  law  of  definite  proportions. — A  pure  substance, 
whether  an  element  or  a  compound,  has  always  the  same 
composition.     It  has  been  pointed  out  that  not  the  slight- 
est variation  in  the  proportions  of  hydrogen  and  oxygen  in 
water  can  be  detected.    This  is  true  of  all  compounds  with- 
out exception,  and  it  is  the  great  distinction  between  them 
and  mixtures.     The  statement  is  commonly  called  the  law 
of  definite  proportions.     We  may  express  it  again  in  this 
way :  "  The  composition  of  a  pure  chemical  compound  is 
constant" 

We  can  now  state  another  difference  between  the  mix- 
ture and  the  compound  of  copper  and  sulphur.  The  mixture 
we  can  make  in  any  desired  proportions  of  the  two.  But  the 
compound  which  we  have  made"  will  always  contain  126 
parts  of  copper  and  32  of  sulphur,  and  we  are  powerless  to 
produce  the  slightest  alteration  in  these  proportions.1 

To  sum  up.  Experiment  reveals  this  remarkable  fact, 
that  there  are  certain  forms  of  matter  whose  composition 
is  constant.  Chemists  agree  to  name  these  substances  chem- 
ical compounds.  When  they  discover  that  the  composition 
of  a  substance  varies,  they  class  it  with  the  mixtures. 

41.  Elements. — In  the  first  chapter  we  separated  water 
into  hydrogen  and  oxygen  by  passing  the  electric  current 
through  it.    Mercuric  oxide  also  we  have  split  up  into  oxy- 
gen and  mercury  by  heat.    By  various  methods  all  chemical 
compounds  can  be  decomposed,  some  by  the  current  like 

1  There  is  another  compound  of  these  two  elements  which  contains 
half  as  much  copper  to  the  same  quantity  of  sulphur — that  is,  63  parts 
copper  and  32  sulphur ;  but  this  is  a  different  substance  and  does  not 
affect  the  argument. 


34  ELEMENTARY  CHEMISTRY 

water,  very  many,  most,  in  fact,  by  intense  heat,  like  mer- 
curic oxide.  As  a  result  of  this  separating  process  we  find 
that  we  have  obtained  substances  which  resist  all  attempts  to 
decompose  them  further.  Chemists  have  never  succeeded  in 
separating  mercury  or  oxygen  or  hydrogen  into  simpler 
forms  of  matter.  £11  of  the  very  numerous  compounds 
which  have  been  found  in  the  earth's  crust  have  been  treated 
in  this  way,  and  as  a  result  there  have  been  obtained  about 
eighty  substances  which,  at  present,  can  not  be  further  de- 
composed. These  are  called  elements.  It  is  impossible  to 
state  the  number  precisely,  because,  owing  to  the  great  rarity 
of  some  of  them,  and  to  the  imperfections  of  our  methods, 
there  are  always  some  whose  elementary  character  is  in  doubt. 
42.  Meaning  of  the  term  "element." — It  is  important  to 
understand  clearly  the  way  in  which  the  idea  of  "  element  " 
is  arrived  at.  An  element  is  simply  a  substance  which,  at 
present,  can  not  be  further  separated.  As  we  progress  we 
shall  discover  examples  of  bodies  once  thought  elements 
which  turned  out  to  be  compounds  when  improved  methods 
of  separation  were  applied  to  them.  Thus  we  are  never  in 
a  position  to  assert  that  any  substance  will. always  resist 
the  efforts  of  chemists  to  decompose  it.  All  we  can  say  is 
that  it  has  thus  far  resisted  those  efforts.  But  if  the  ele- 
ments are  really  compounds  of  hitherto  unknown  sub- 
stances, then  these  substances  are  knit  together  far  more 
firmly  than  the  constituents  of  any  ordinary  compound. 
This  is  shown  not  only  by  the  fact  that  they  resist  all  at- 
tempts to  split  them  up  in  the  laboratory,  but  also  by  the 
work  of  astronomers,  who  have  detected  many  of  them  in 
the  atmospheres  of  the  sun  and  other  stars,  where  they 
resist  temperatures  far  higher  than  those  we  can  produce 
artificially.1 

1  Recent  results  make  the  temperature  of  the  sun  about  6,000°  C. 
Many  stars  are  far  hotter  than  this.  The  highest  temperature  we  can 
produce  in  the  laboratory  does  not  exceed  3,500°  C. 


CHAPTER  V 

HYDROGEN 

43.  Occurrence. — Hydrogen  compounds  exist  in  the 
bodies  of  all  animals  and  plants,  and  many  of  them  are 
found  in  the  rocks  as  minerals. 

The  gas  itself  is  rare  in  nature.  It  has  been  found  en- 
closed in  meteors  and  in  masses  of  rock-salt,  it  is  contained 
in  the  natural  gas  of  Pennsylvania  and  in  the  intestinal 
gases  of  animals,  and  is  a  product  of  the  complex  chemical 
process  which  we  call  putrefaction  or  decay.  Hydrogen 
is  continually  being  thrown  into  the  atmosphere  by  the 
gases  from  volcanoes,  and  just  what  becomes  of  this  hydro- 
gen is  a  matter  of  great  interest.  We  should  expect  it 
to  accumulate,  but  it  does  not,  for  although  even  the  pure 
air  over  the  ocean  contains  a  little  of  the  gas,  the  quantity 
is  very  small.  Its  extreme  lightness  makes  it  appear  prob- 
able that  it  escapes  continually  from  the  upper  regions 
of  the  air  into  space,  and  so  its  accumulation  is  prevented.1 

•44.  Preparation  from  sodium  and  water. — Two  methods 
of  obtaining  hydrogen  from  water  have  already  been  men- 
tioned, the  passage  of  the  electric  current  and  the  applica- 
tion of  intense  heat.  Sodium  is  a  soft,  wax-like  metal,  which 
must  be  kept  under  naphtha,  as  it  is  rapidly  acted  upon  by 
air.  A  cylinder  is  filled  with  water  and  inverted  in  a  dish 

1  Enormous  quantities  of  free  hydrogen  are  contained  in  the 
atmosphere  of  the  sun  and  other  stars.  These  stars  are  immensely 
larger  than  the  earth,  and  gravitation  is  sufficiently  intense  to  prevent 
the  hydrogen  from  escaping. 

4  35 


36 


ELEMENTARY  CHEMISTRY 


FIG.  8. — Hydrogen  from  water  and  sodium. 


of  water,  and  a  small  piece  of  sodium,  which  must  be  per- 
fectly clean,  is  wrapped  tightly  in  wire  gauze,  and  the  gauze, 

held  in  forceps,  is 
introduced  under 
the  mouth  of  the 
tube  (Fig.  8). 
Bubbles  of  a  color- 
less gas  rise  into 
the  tube.  The  so- 
dium has  disap- 
peared and  the  wa- 
ter in  the  dish  has 
acquired  a  soapy 
feel.  This  is  due  to 
the  presence  of  so- 
dium hydroxide, 
commonly  called  caustic  soda,  which  is  dissolved  in  'the  water, 
and  which  we  could  obtain  as  a  white  solid  by  evaporating 
the  solution  to  dryness.  Caustic  soda  is  a  compound  of  sodi- 
um, oxygen,  and  hydrogen.  The  sodium  has  combined  with 
all  of  the  oxygen  and  half  of  the  hydrogen  of 
the  water,  and  the  other  half  of  the  hydrogen  /  ••* .fc,  , 
has  been  liberated.  Of  course,  only  a  small 
portion  of  the  water  has  taken  part  in  the 
change.  The  rest  remains,  and  in  it  the  sodi- 
um hydroxide  is  dissolved. 

We  can  easily  prove  that  the  cylinder  con- 
tains a  combustible  gas  by  taking  it  out  of  the 
water  and  bringing  the  open  end  near  a  burn-  Fm-  9--Hydr°- 

gen    from    hy- 
ing match.    A  pale  flame  begins  at  the  mouth      drochioric  acid 

of  the  cylinder  and  runs  to  the  bottom.     The      and  zinc- 
hydrogen  combines  with  oxygen  from  the  air,  and  again 
exists  as  water. 

45.  From  zinc  and  acids. — Some  zinc  is  placed  in  a  beak- 
er, and  hydrochloric  acid  poured  upon  it.    There  is  a  violent 


HYDROGEN  37 

escape  of  gas,  which,  on  being  lighted,  burns  with  a  large 
pale  flame  (Fig.  9).  The  gas  is  hydrogen.  Hydrochloric 
acid  is  a  compound  of  hydrogen  and  chlorine  and  the  zinc 
combines  with  the  chlorine,  producing  zinc  chloride,  while 
the  hydrogen  escapes. 

For  preparing  larger  quantities  of  hydrogen,  the  appa- 
ratus shown  in  Fig.  10  is  employed.  It  consists  of  a  large 
bottle,  closed  by  a 
doubly  perforated  rub- 
ber cork.  Through 
one  hole  runs  a  tube 
which  has  a  small  fun- 
nel at  the  upper  end. 
The  lower  end  of  this 
tube  reaches  nearly  to 
the  bottom  of  the  bot- 
tle. The  other  hole 

•    i  1-1  FIG.  10. — Preparation  of  hydrogen. 

carries   a   tube   which 

ends  just  inside  the  cork,  and  through  it  the  gas  is  led 
away.  In  the  bottle  is  placed  some  zinc  covered  with  water, 
and  sulphuric  acid  is  poured  through  the  funnel  tube.  At 
once  the  evolution  of  hydrogen  begins.  This  first  quantity 
of  gas  mixes  with  the  air  in  the  bottle  and  produces  an  ex- 
plosive mixture,  so  that  it  is  necessary  to  wait  until  all  the 
air  is  expelled  from  the  apparatus  before  collecting  or 
lighting  the  gas.  The  gas  may  be  collected  over  water,  as 
shown  in  the  drawing. 

46.  Physical  properties. — Hydrogen  is  a  colorless,  odor- 
less gas.  By  cooling  the  strongly  compressed  gas  to  an  ex- 
tremely low  temperature,  and  then  permitting  it  to  escape 
through  a  small  orifice  into  an  intensely  cold  vessel,  it  has 
been  converted  into  a  colorless  liquid.  When  still  more 
intensely  cooled,  it  solidifies  to  an  ice-like  solid.  The  evap- 
oration of  liquid  hydrogen  produces  the  lowest  tempera- 
ture known  to  science,  a  temperature  only  about  13°  above 


38  ELEMENTARY  CHEMISTRY 

the  absolute  zero  of  physics.  When  an  empty  test-tube  is 
dipped  into  it,  solid  air  at  once  collects  in  the  tube.  Liquid 
hydrogen  is  less  than  one-tenth  as  dense  as  water,  and  solid 
hydrogen  appears  to  have  about  the  same  density  as  the 
liquid.  The  liquid  is  therefore  by  far  the  lightest  of  all 
liquids,  and  the  solid  of  all  solids.  The  gas  is  the  lightest 
of  gases,  air  being  14.39  times  as  heavy.  It  is  very  slightly 
soluble  in  water,  50  volumes  of  water  dissolving  1  volume 
of  the  gas,  at  ordinary  temperatures.  It  has  been  seen  that 
hydrogen  is  extremely  difficult  to  liquefy,  and  it  is  an  in- 
teresting fact  that  this  is  the  case  with  all  gases  whose 
solubility  is  very  slight. 

Hydrogen  dissolves  in  various  metals,  in  gold,  silver, 
and  platinum,  for  instance,  and  notably  in  palladium,  a 
white  metal  somewhat  similar  to  platinum.  Palladium  will 
dissolve  about  1,000  times  its  volume  of  the  gas.  This  power 
of  hydrogen,  of  forming  solid  solutions  with  the  metals, 
explains  its  ability  to  pass  through  thin  plates  of  iron  or 
platinum. 

47.  Chemical  properties. — When  a  current  of  hydrogen 
is  ignited  where  it  escapes  into  the  air  it  burns  with  a  pale 
blue  flame,  combining  with  the 
oxygen  of  the  air  to  produce  wa- 
ter.1 It  has  been  shown  that  some 
little  hydrogen  escapes  from  the 
flame  into  the  air  without  being 
burned.  When  hydrogen  burns  in 
oxygen  it  produces  the  hottest 

Fia.  11. — Oxy hydrogen  blowpipe,     n  V»*   V,     V,          V>  Ivf    '       r\ 

the  oxyhydrogen  flame.  This  melts  platinum  easily,  and 
is  used  in  working  the  metal  into  crucibles  and  dishes  for 

1  It  is  best  to  insert  a  short  tube  made  of  platinum  foil  into  the 
glass  tube,  otherwise  the  flame  is  colored  yellow  by  sodium  from  the 
glass. 


HYDROGEN 


39 


chemical  purposes.     Solids — lime,  for  instance — which  do 
not  melt  when  placed  in  the  flame,  glow  with  a  bright 


FIG.  12.— Burning  of  iron  in  the  flame  of  th 


n  blowpipe. 


light,  the  so-called  lime-light  or  Drummond  light.  The 
apparatus  usually  employed  for  producing  the  oxyhydrogen 
flame  on  the  lecture-table  is  represented 
in  Figs.  11  and  12. 

A  cylinder,  filled  with  hydrogen,  is 
held  open  end  down,  and  a  lighted  can- 
dle stuck  on  a  stiff  wire  is  pushed  up  into  it 
(Fig.  13).  The  hydrogen  is  inflamed  at 
the  mouth  of  the  cylinder,  but  the  candle 
goes  out  in  the  gas.  The  burning  of  the 
candle  is  the  combination  of  its  constitu- 
ents with  the  oxygen  of  the  air,  and,  of  PIG  18._Behavior  of 
course,  this  can  not  continue  when  the  can- 
dle is  surrounded  by  hydrogen.  Hydrogen 
is  not  poisonous,  but  small  animals  placed  in  it  suffocate, 
owing  to  the  lack  of  oxygen,  without  which  no  animal  life 
can  continue, 


a  lighted    candle    in 
hydrogen. 


CHAPTER  VI 

OXYGEN  AND  HYDROGEN  PEROXIDE 

48.  The  free  oxygen  of  the  air. — We  may  now  study 
the  other  constituent  of  water.  We  have  already  obtained 
oxygen  by  the  decomposition  of  water,  and  by  heating  mer- 
curic oxide,  as  a  colorless  gas  which  causes  a  spark  to  burst 
into  flame,  and  it  has  been  pointed  out  that  this  is  due 
to  the  fact  that  the  air  contains  only  about  21  per  cent 
by  volume  of  oxygen,  and  that  the  other  constituents  inter- 
fere with  combustion.  This  free  oxygen  of  the  air  is  essen- 
tial to  all  combustion  processes,  *  and  to  the  life  of  all 
animals  and  most  plants.  Under  the  influence  of  sunlight 
the  carbon  dioxide  of  the  air  is  decomposed  in  the  green 
parts  of  plants,  the  carbon  built  up  into  the  structure  of 
the  plant,  and  the  oxygen  returned  to  the  air.  If  some 
sprigs  of  mint  are  placed  in  a  flask  which  is  filled  with 
water  containing  carbon  dioxide,  then  inverted  and  placed 
in  sunlight,  bubbles  of  oxygen  collect  in  the  upper  part  of 
the  flask.  But  the  roots  and,  in  general,  the  portions  of 
plants  which  are  not  green  absorb  oxygen  from  the  air,  and 
give  off  carbon  dioxide.  Thus,  in  a  plant,  two  opposite  pro- 
cesses occur  together,  one  of  which  puts  oxygen  into  the 
air  and  the  other  takes  it  out;  but  the  first  takes  place  on 
a  far  larger  scale,  and  a  plant,  on  the  whole,  absorbs  carbon 
dioxide  from  the  air  and  gives  off  oxygen.  On  the  other 
hand,  combustion  processes  and  the  life  of  animals  make 
continual  drafts  upon  the  oxygen  of  the  air,  and  discharge 
carbon  dioxide  into  it. 
40 


OXYGEN  AND  HYDROGEN  PEROXIDE  41 

If  these  two  opposed  influences  were  perfectly  balanced, 
the  percentages  of  oxygen  and  of  carbon  dioxide  in  the  air 
would  remain  the  same.  Only  slight  variations  with  the 
seasons  have  been  detected.  Air  in  a  vase  discovered  in  the 
ruins  of  Pompeii,  which  had  been  sealed  up  for  nearly  two 
thousand  years,  was  analyzed  by  Liebig,  and  was  found  to 
possess  the  same  composition  as  the  air  which  bathes  the  sur- 
face of  the  planet  at  present.  But  it  would  take  a  long  time 
to  produce  a  variation  which  could  be  detected,  and  we  are 
by  no  means  justified  in  assuming  that  the  two  processes  in 
question  exactly  compensate  each  other,  more  especially  as 
we  know  from  geological  facts  that  there  was  a  period  in  the 
history  of  the  earth  when  the  plants  obtained  the  upper 
hand,  and  produced  an  enormous  decrease  in  the  carbon 
dioxide  and  increase  in  the  oxygen  of  the  air.  At  present  it 
is  very  probable  that  the  pendulum  has  swung  in  the  other 
direction,  and  that  the  immense  combustion  processes  car- 
ried out  by  man  are  gradually  increasing  the  carbon  dioxide 
and  decreasing  the  oxygen. 

In  spite  of  its  importance,  the  quantity  of  oxygen  in  the 
air,  about  21  per  cent  by  volume  and  23  per  cent  by  weight, 
is  small  compared  with  that  which  exists  in  oxygen  com- 
pounds. Water  contains  eight-ninths  of  its  weight  of  oxy- 
gen, and  a  layer  of  water  264  centimeters  (8|  feet)  in 
thickness  over  the  earth  would  contain  as  much  oxygen  as 
the  atmosphere. 

49.  Preparation. — Potassium  chlorate  is  a  white  crys- 
talline compound  of  potassium,  chlorine,  and  oxygen.  When 
it  is  heated  it  first  melts,  and  then  gives  off  its  oxygen,  leav- 
ing potassium  chloride,  a  compound  of  potassium  and  chlo- 
rine, behind.  Since  the  potassium  chlorate  may  explode 
when  heated  alone,  it  is  usual  to  mix  with  it  some  substance 
which  causes  the  oxygen  to  be  given  off  more  regularly,  and 
at  a  lower  temperature.  Various  substances  possess  this 
power.  Finely  divided  platinum  is  one,  but  the  material 


42  ELEMENTARY  CHEMISTRY 

commonly  employed  is  manganese  dioxide.  The  manganese 
dioxide  is  found  unaltered  after  the  experiment,  mixed  with 
the  potassium  chloride.  It  promotes  the  decomposition  of 
the  potassium  chlorate,  but  remains  unaffected  itself.  This 
is  called  a  case  of  catalytic  action,,  and  this  term  is  applied 
to  all  cases  in  which  substances  influence  chemical  changes 
into  which  they  do  not  enter. 

50.  Physical  properties. — The  oxygen  can  be  collected 
over  water  in  which  it  is  only  slightly  soluble.    The  appa- 
ratus is  the  same  as  that  shown  in  Fig.  28. 

Oxygen  is  a  colorless,  odorless  gas.  By  great  cold  and 
pressure  it  is  converted  into  a  light  blue  liquid,  which  is 
somewhat  denser  than  water,  and  is  strongly  attracted  by 
the  magnet.  One  hundred  volumes  of  water  at  0°  and  under 
one  atmosphere  pressure  dissolve  less  than  5  volumes  of 
the  gas,  so  that  the  solubility  is  slight.  All  natural  sur- 
face waters  contain  dissolved  oxygen,  and  while  the  quan- 
tity is  small  it  is  important,  for  .without  it  fish-life  would 
be  impossible.  The  dissolved  oxygen  can  be  extracted  from 
water  by  boiling  and  cooling  in  the  absence  of  air,  or  by 
means  of  the  air-pump.  In  such  water  fish  suffocate. 

51.  Chemical  properties. — The  combustions  which  occur 
in  the  air  consist  in  the  combination  of  carbon  and  hydro- 
gen with  the  oxygen  of  the  atmosphere.    These  phenomena 
occur  far  more  energetically  in  pure  oxygen  than  in  air.    At 
present  oxygen  is  no  longer  made  in  large  quantities  on  the 
chemical  lecture-table,  because  the  steel  cylinders,  full  of 
the  compressed  gas,  which  are  sold  furnish  a  far  more  con- 
venient source.    From  one  of  these  we  collect  over  water  a 
number  of  jars  full  of  oxygen.    A  little  sulphur  is  placed  in 
an  iron  spoon  and  heated  in  the  air  until  it  begins  to  burn 
feebly.    When  plunged  into  oxygen  there  is  a  great  increase 
in  the  vigor  of  the  combustion  (Fig.  14).     Charcoal,  espe- 
cially that  made  from  bark,  burns  in  it,  throwing  off  show- 
ers of  sparks.    A  steel  watch-spring — steel  we  can  consider 


OXYGEN  AND  HYDROGEN   PEROXIDE  43 

as  a  variety  of  iron — whose  combustion  has  been  started  by 
a  piece  of  wood  split  and  slipped  over  the  end,  is  converted 
into  iron  oxide,  with  brilliant  scintillations  (Fig.  15). 
Phosphorus  and  magnesium  burn  with  blinding  splendor, 
and  with  such  intense  evolution  of 
heat  that  the  jar  is  often  frac- 
tured. In  all  these  cases  the  prod- 
ucts are  the  oxides  of  the  ele- 
ments employed,  those  of  carbon 
and  sulphur  being  colorless  gases 
and  those  of  phosphorus  and  mag- 
nesium white  solids.  In  absolute- 
ly dry  oxygen,  charcoal  and  phos- 
phorus can  scarcely  be  induced  to  Fl°-  ^.-combustion  of  sui- 

J  -phur  in  oxygen. 

burn   at   all,   and   there   are   many 

other  combustions  which  only  take  place  in  the  presence  of 
traces  of  water-vapor.  We  can  now  classify  this  action  of 
water  with  the  action  of  manganese  dioxide  on  potassium 
chlorate  as  a  case  of  catalytic  action. 

52.  Relation  of  oxygen  to  life. — Free  oxygen  is  essential 
to  the  life  of  animals.  In  its  absence  suffocation  results. 
Thus,  when  a  small  animal,  like  a 
mouse,  is  plunged  into  a  jar  of  ni- 
trogen, it  dies  at  once;  though  the 
nitrogen  is  not  at  all  poisonous,  it 
kills  by  excluding  the  oxygen,  with- 
out which  life  can  not  continue. 
The  energy — heat  and  motion— 
which  continually  appears  in  the 
animal  body  during  life  has  its 
FIG.  15.— combustion  of  iron  S0urce  in  the  combination  of  the 
carbon  and  hydrogen  of  the  tissues 

with  the  oxygen  of  the  air.  The  products  of  this  change  are 
carbon  dioxide  and  water,  just  as  in  combustions  which  take 
place  outside  the  body.  For  this  reason  the  air  which 


44  ELEMENTARY  CHEMISTRY 

leaves  the  lungs  has  a  very  different  composition  from  that 
which  enters  them.  It  contains  more  than  4  per  cent  by 
volume  of  carbon  dioxide — about  one  hundred  times  as 
much  as  the  inspired  air — and  has  lost  about  5  per  cent  of 
the  oxygen. 

Oxygen.    Nitrogen  and  argon.    Carbon  dioxide. 

Pure  air 20.96  79.01  0.03         )  Per  cent  by 

Expired  air 16.03  79.59  4.38          )      volume. 

Carbon  dioxide  is  eliminated  and  oxygen  absorbed 
through  the  skin  also,  but  the  quantities  are  trifling. 

Oxygen  is  essential  to  the  higher  plants,  though  death, 
when  it  is  withdrawn,  is  far  less  rapid  than  with  animals. 
It  is  an  interesting  fact  that  although  oxygen  is  the  con- 
stituent of  the  air  which  is  most  directly  concerned  with 
the  life  process,  it  is  not  able  to  support  life  alone.  An  ani- 
mal supplied  with  the  pure  gas  constantly  renewed  will  die 
at  length,  though  life  continues  |or  a  long  time. 

OZONE 

53.  In  the  experiment  with  mercuric   iodide  we   have 
met  with  a  substance  which  can  exist  in  two  very  different 
modifications.     It  will  not  surprise  us,  therefore,  to  learn 
that  many  of  the  elements  possess  the  same  power.    Oxygen 
furnishes  an  interesting  example  of  this. 

54.  Preparation. — For  more  than  a  century  it  has  been 
known  that  the  air  near  an  electrical  machine  in  operation 
acquires  a  peculiar  odor.     This  is  due  to  the  fact  that  the 
oxygen  of  the  air  is  transformed  by  electric  discharges  into 
another  modification,  ozone,  which  has  a  powerful  smell. 
In  order  to  prove  this,  it  is  best  to  employ  pure  oxygen, 
since  otherwise  it  might  reasonably  be  thought  that  the 
nitrogen,  or  some  of  the  other  constituents  of  the  air,  takes 
some  part  in  the  change. 

Some  starch  paste  is  prepared  by  boiling  a  little  starch 
with  water  in  a  dish.  Into  a  small  portion  of  this  in  a 


OXYGEN  AND  HYDROGEN  PEROXIDE  45 

test-tube  throw  a  little  powdered  iodine.  An  intense  blue 
color  is  produced,  and  this  color  always  results  when  free 
iodine  and  starch  paste  come  into  contact.  To  the  rest  of  the 
liquid  a  crystal  of  potassium  iodide  is  added,  and  it  remains 
colorless.  Notice  that  iodine  in  the  form  of  an  iodide  pro- 
duces no  color  with  starch  paste,  but,  of  course,  anything 
which  sets  free  the  iodine  will  immediately  produce  the  blue 
tint  in  the  mixture.  A  long  glass  tube,  a,  sealed  at  one  end, 
is  filled  with  this  liquid  and  inverted  in  more  of  it  in  a 
beaker,  &,  and  some  oxygen  is  passed  up  into  the  tube.  There 
is  no  coloration,  oxygen  not  having  any  action  on  potassium 


FIG.  16.— Transformation  of  oxygen  into  ozone  by  electric  sparks. 

iodide.    This,  of  course,  we  might  have  predicted  from  the 
fact  that  the  liquid  remained  colorless  in  the  air. 

Through  the  walls  of  the  tube  near  the  top  there  pass 
two  platinum  wires  which  enter  opposite  each  other  and  do 
not  quite  meet  in  the  center.  These  wires  are  now  connected 
with  an  induction  coil  and  a  stream  of  sparks  passed  through 
the  oxygen  (Fig.  16).  Gradually  the  blue  color  makes  its 
appearance  at  the  top  of  the  liquid  and  spreads  slowly  down- 
ward. The  sparks  have  transformed  a  portion  of  the 
oxygen  into  ozone,  which  attacks  the  potassium  iodide  and 


46  ELEMENTARY  CHEMISTRY 

liberates  the  iodine.  The  latter  then  produces  the  blue 
color  with  the  starch.  Far  better  than  sparks  for  the  pro- 
duction of  ozone  is  the  silent  electric  discharge.  A  wide 
tube  is  coated  outside  with  tin-foil,  and  a  narrow  tube,  which 


FIG.  17.— Preparation  of  ozone. 

• 

fits  in  it,  is  coated  inside  in  the  same  way.  Through  the 
ring-shaped  space  between  the  two  tubes  oxygen  is  passed 
slowly  from  d  to  c,  the  coatings  being  connected  at  a  and  b 
with  the  induction  coil  (Fig.  17).  The  issuing  gas  con- 
tains large  quantities  of  ozone,  as  well  as  much  unchanged 
oxygen,  from  which  it  is  best  separated  by  intense  cooling, 
when  ozone  containing  only  a  little  oxygen  condenses.  If 
the  liquid  is  then  allowed  to  warm  a  little,  the  oxygen  vapor- 
izes first  and  nearly  pure  ozone  is  obtained. 

This  method  of  preparing  ozone  is  very  important,  be- 
cause we  have  made  it  from  oxygen  without  adding  or  sub- 
tracting anything  except  energy,  and  there  is  no  escape 
from  the  conclusion  that  it  is  a  variety  of  oxygen.  There 
are  many  other  methods  of  making  it,  but  in  all  it  is  ob- 
tained mixed  with  large  quantities  of  ordinary  oxygen. 
Thus,  when  a  stick  of  phosphorus  is  placed  in  a  jar  and 
partly  covered  with  water,  the  air  in  the  jar  acquires  the 
odor  of  ozone  and  will  turn  blue  a  paper  which  has  been 
dipped  in  the  solution  of  potassium  iodide  and  starch  (Fig. 


OXYGEN  AND  HYDROGEN    PEROXIDE  47 

18).     When  barium  peroxide  is  acted  upon  by  strong  sul- 
phuric acid,  oxygen  quite  rich  in  ozone  is  given  off,  and  can 
be  identified  by  the  odor,  by  the  starch  and  potassium-iodide 
test  and  by  its  action  on  silver.     It  is 
sufficient  to  cover  the  test-tube  with  a 
clean  silver  coin  when  a  dark  stain  of 
silver  peroxide  is  produced  upon  it. 

55.  Properties. — Ozone  is  a  color- 
less gas   with   a   peculiar   penetrating 
odor — the  odor  of  phosphorus  is  due  to 
it.1  It  is  just  1J  times  as  dense  as  oxy- 
gen.    The  density  of  oxygen  referred 
to    hydrogen    is    nearly    16;    that    of 

ozone  nearly  24.  Thus,  when  oxygen  FIG.  is.— Ozone  from  phos- 
is  converted  into  ozone,  1J  liters  is 
crowded  into  the  space  of  one.  liter,  or  three  liters  of  oxygen 
yield  two  of  ozone.  It  is  liquefied  by  cold  and  pressure 
more  easily  than  oxygen,  and  the  liquid  is  blue-black,  mag- 
netic, and  extremely  explosive.  The  explosion  is  the  sudden 
conversion  of  the  ozone  into  ordinary  oxygen,  a  process  in 
which  energy  is  set  free  and  increase  of  volume  occurs. 

56.  Change  of  ozone  to  ordinary  oxygen. — This  change 
of  ozone  to  oxygen  occurs  slowly  at  ordinary  temperatures. 
If  some  ozone  is  sealed  up  in  a  glass  bulb  it  is  found,  after 
a   time,   that   the   bulb   contains   ordinary   oxygen.     Heat 
greatly  quickens  the  change.     If  a  stream  of  oxygen  con- 
taining ozone  is  passed  through  a  glass  tube  gently  heated 
by  brushing  it  with  a  flame,  the  gas  which  issues  will  be 
ordinary  oxygen  and  will  produce  no  color  upon  the  potas- 
sium-iodide starch-paper.     On  the  other  hand,  at  extremely 
high   temperatures,   oxygen   changes    to    ozone.      A    large 
flask  is  filled  with  oxygen  over  water  and  a  flame  of  hydro- 
gen, led  from  a  generator  and  burning  at  a  platinum  tip, 

1  In  thick  layers  ozone  is  blue. 


48  ELEMENTARY  CHEMISTRY 

is  introduced.  After  a  moment  the  flame  is  withdrawn, 
some  of  the  solution  of  potassium  iodide  containing  starch 
introduced,  the  flask  closed  and  vigorously  shaken.  The 
blue  color  appears  at  once.  The  high  temperature  of  the 
hydrogen  flame  has  converted  some  of  the  oxygen  into  ozone. 

57.  Chemical  properties. — Chemically,  ozone  has  all  the 
properties  of  oxygen.     The  combustions  of  charcoal,  sul- 
phur, phosphorus,  and  so  on,  which  we  have  carried  out  in 
oxygen,  will  occur  in  ozone  even  more  energetically,  because 
the  amount  of  heat  evolved  is  greater.     But  ozone  is  far 
more  active  chemically  than  ordinary  oxygen.     It  attacks 
and  destroys  organic  matter  like  paper  and  rubber;  blood 
becomes  colorless  in  contact  with  it.     Silver,  which  is  not 
acted  upon  by  ordinary  oxygen  at  any  temperature,  is  con- 
verted into  black  silver  peroxide,  and  the  production  of  this 
black  stain  on  a  polished  silver  surface  is  the  most  certain 
evidence  of  the  presence  of  ozone  *we  possess.1    The  reputa- 
tion which  ozone  has  obtained  for  health-giving  qualities  is 
entirely  undeserved.     In  any  considerable  quantity  the  gas 
is  violently  poisonous,  and  even  when  very  much  diluted 
with  air  it  is  irritating  and  dangerous. 

HYDROGEN  PEROXIDE 

58.  Preparation. — Barium  peroxide  is  a  white  compound 
of  barium  and  oxygen.     It  has  been  remarked  that  when 
strong  sulphuric  acid  acts  upon  it,  oxygen  containing  ozone 
escapes.    We  have  noted  that  sulphuric  acid  is  a  compound 
of  hydrogen,  sulphur,  and  oxygen.    When  it  comes  into  con- 
tact with  barium  peroxide,  the  sulphur  and  the  oxygen  of 
the  acid  combine  with  the  barium,  producing  a  white  solid 
called  barium  sulphate,  which  remains  in  the  tube,  the 
hydrogen  of  the  acid  unites  with  half  of  the  oxygen  of  the 
peroxide,  producing  water,  while  the   other  half  of  the 

1  Unfortunately  this  test  is  far  from  delicate. 


OXYGEN  AND  HYDROGEN  PEROXIDE  49 

oxygen  escapes,  partly  as  ozone.  Now,  when  barium  per- 
oxide is  slowly  added  to  cold  dilute  sulphuric  acid,  barium 
sulphate  is  produced  in  the  same  way,  but  no  gas  escapes. 
In  this  case  the  hydrogen  of  the  acid  unites  with  all  of  the 
oxygen  of  the  barium  peroxide,  producing  a  compound  which 
must  contain  twice  as  much  oxygen  for  a  given  weight  of 
hydrogen  as  water,  and  which  is  called  hydrogen  peroxide. 
The  barium  sulphate,  being  insoluble  in  water,  can  be  easily 
got  rid  of  by  filtering  the  liquid,  and  the  clear  liquid  which 
passes  through  is  a  solution  of  hydrogen  peroxide  in  water. 

59.  Properties. — Some  of  this  liquid  is  allowed  to  fall 
on  a  little  manganese  dioxide  in  a  test-tube.  There  is  a  brisk 
evolution  of  gas,  which  is  found  on  trial  to  respond  to  the 
spark  test.  Even  when  alone,  hydrogen  peroxide  gives  up 
half  its  oxygen  passing  into  water.  This  decomposition  into 
oxygen  and  water  is  very  slow  when  it  is  preserved  in  the 
dark  and  somewhat  more  rapid  in  the  light.  It  becomes  very 
rapid  when  the  hydrogen  peroxide  is  heated  or  brought  into 
contact  with  manganese  dioxide.  Finely  divided  platinum, 
gold,  silver,  or  mercury  produce  the  same  effect.  These  are 
further  cases  of  what  we  have  called  "  catalytic  action," 
and  it  is  interesting  to  notice  that,  in  this  case,  the  effect 
of  the  active  substance  is  to  increase  very  much  the  speed 
of  a  chemical  change  which  would  nevertheless  occur  alone 
if  time  enough  were  given.  There  is  excellent  reason  to 
believe  that  this  is  always  the  case,  and  that  no  substance 
acting  in  this  way  can  alter  the  final  state  of  things,  but 
only  the  speed  with  which  that  final  state  is  reached. 

We  now  pour  some  of  our  hydrogen  peroxide  on  a  little 
lead  sulphide,  a  black  compound  of  lead  and  sulphur.  Lead 
sulphate,  a  white  compound  of  lead,  sulphur,  and  oxygen,  is 
produced.  Most  paints  contain  lead  compounds,  and  the 
blackening  of  oil-paintings  with  age  is  due  to  the  forma- 
tion of  lead  sulphide  by  the  action  of  sulphur  gases  in  the 
atmosphere.  Careful  treatment  with  hydrogen  peroxide 


50  ELEMENTARY   CHEMISTRY 

converts  the  dark  lead  sulphide  into  white  lead  sulphate  and 
restores,  in  a  measure,  the  original  tints. 

From  all  this  it  is  clear  that  the  additional  quantity 
of  oxygen  in  hydrogen  peroxide  is  only  loosely  held;  it 
easily  yields  it  up  to  any  oxidizable  substance,  passing  into 
water  at  the  same  time,  and  thus  it  communicates  oxygen 
to  other  substances,  or  acts  as  an  oxidizing  agent.  Never- 
theless, under  some  circumstances,  it  may  remove  oxygen, 
or  produce  reduction.  Some  silver  oxide  is  placed  in  a 
beaker  and  hydrogen  peroxide  poured  over  it.  There  is  a 
brisk  evolution  of  oxygen,  and  metallic  silver  and  water 
remain.  In  this  case  the  loosely  held  oxygen  of  the  silver 
oxide  and  of  the  hydrogen  peroxide  escape  together.  Hy- 
drogen peroxide  acts  in  a  similar  way  with  the  oxides  of 
gold  and  mercury. 

The  purest  hydrogen  peroxide  is  a  thick  colorless  liquid, 
perfectly  soluble  in  water,  alcohol,  and  ether.  Even  in  the 
cold,  and  rapidly  when  heated,  it  decomposes  into  water 
and  oxygen,  and,  under  the  influence  of  heat,  the  decompo- 
sition frequently  occurs  with  violent  explosion.  When 
water  evaporates  in  the  air  a  small  quantity  of  it  com- 
bines with  the  oxygen  of  the  atmosphere,  producing  hydro- 
gen peroxide.  For  this  reason,  traces  of  hydrogen  peroxide 
are  always  contained  in  the  atmosphere,  and  in  rain-water 
and  snow. 

The  water  solution  of  hydrogen  peroxide  is  a  colorless 
transparent  liquid.  It  is  largely  used  for  bleaching,  and  in 
surgery  for  washing  wounds. 


CHAPTER   VII 

COMBUSTION 

60.  SULPHUR,  and  especially  phosphorus,  combine  slow- 
ly with  oxygen  at  ordinary  temperatures.  In  the  case  of 
sulphur  no  evolution  of  light  or  heat  can  be  detected.  Phos- 
phorus is  luminous  in  the  dark  owing  to  this  gradual 
oxidation.  When  either  substance  is  heated,  the  combina- 
tion becomes  more  and  more  rapid.  This  is  usually  the 
effect  of  raising  the  temperature;  to  increase  enormously 
the  speed  with  which  a  chemical  change  takes  place. 
Finally,  if  the  burning  substance  is  plunged  into  oxygen, 
there  is  a  still  further  increase  in  the  rapidity  of  the  pro- 
cess. Now  in  all  these  cases  the  total  amount  of  heat  pro- 
duced by  the  combination  of  a  given  weight  with  oxygen 
is  exactly  the  same.  In  the  same  way,  when  a  metal  rusts, 
slowly  forming  its  oxide  in  a  process  requiring  years,  the 
same  amount  of  heat  is  produced  as  though  it  burned  in 
oxygen  in  a  few  seconds,  provided  only  that  the  product 
is  the  same  in  both  cases.  The  only  difference  is  that,  in 
the  first  case,  the  evolution  of  heat  is  very  slow,  it  is  dissi- 
pated as  fast  as  produced,  and  there  is  no  perceptible  eleva- 
tion of  temperature.  But,  in  the  second  case,  the  energy 
is  evolved  in  a  few  seconds,  the  substances  are  intensely 
heated,  and  the  striking  phenomena  of  combustion  appear. 
The  key  to  all  this  is  the  increase  in  the  speed  of  the  change 
with  rising  temperature.  For,  as  soon  as  the  change  pro- 
duces heat  too  rapidly  for  it  to  be  dissipated,  this  heat  will 
accumulate  and  raise  the  temperature  of  the  substances. 
5  51 


52 


ELEMENTARY  CHEMISTRY 


This  will  produce  an  increase  in  the  speed  of  the  change, 
which  will  in  turn  produce  more  heat  and  a  higher  tem- 
perature, and  so  on. 

61.  Ordinary  combustion  consists  in  the  rapid  combina- 
tion of  various  substances  with  the  oxygen  of  the  air.  Such 
substances  are  called  combustible.  A  gas,  like  oxygen, 
which  permits  them  to  burn,  is  said  to  support  combustion ; 
and  this  language  is  applied  to  any  gas  which  permits  the 
same  substances  to  burn  in  it.  Chlorine  gas — one  of  the 
constituents  of  common  salt — combines  with  phosphorus 
and  other  substances,  producing  light  and  heat,  and  is 
usually  called  a  supporter  of  combustion.  It  is  very  de- 
sirable to  perceive  that  whether  a  substance  will  burn  or 
not  depends  entirely  on  what  particular  gas  surrounds  the 
substance  at  the  time  the  experiment  is  made;  and  that 

whether  a  gas  "  supports  com- 
bustion "  or  not  depends  upon 
the  substance  we  try  to  burn  in 
it.  We  have  seen  that  a  candle 
becomes  incombustible  in  an 
atmosphere  of  hydrogen  (p. 
39). 

A  convenient  apparatus  for 
experiments     on    this     subject 
consists   of  an  ordinary  lamp- 
chimney  closed  below  by  a  rub- 
ber   cork.      Through   the    cork 
passes  a  glass  tube,  by  means 
FIG.  19.— Apparatus  for  combustion    of  which  a   current   of  hydro- 
gen 1  is  introduced  from  a  gen- 
erator or  from  a  cylinder  of  the  compressed  gas    (Fig. 
19).     The  open  top  of  the  chimney  is  covered  with  a  piece 

1  The  use  of  a  hydrogen  generator  can  be  avoided  by  using  common 
illuminating  gas  directly  from  the  mains.  This  serves  the  same  pur- 
pose as  hydrogen  in  these  experiments. 


COMBUSTION  53 

of  asbestos  cardboard,  a,  having  a  hole  in  the  center  through 
which  the  gas  issues,  and  where  it  can  be  lighted  to  avoid 
its  escape  into  the  air.  We  have,  then,  in  the  chimney 
an  atmosphere  of  hydrogen,  and  we  can  conveniently  in- 
vestigate the  behavior  of  various  substances  by  heating 
them  in  iron  deflagrating  spoons  and  introducing  them 
into  the  hydrogen  through  the  hole  in  the  asbestos  plate. 

62.  Combustions  in  hydrogen. — First,  let  some  turpen- 
tine be  started  burning  in  the  air  and  then  plunged  into  the 
hydrogen.     The  flame  is  extinguished.     Like  the  candle, 
turpentine  is  composed  of  carbon  and  hydrogen  and  can 
not  continue  burning  in  a  gas  which  does  not  supply  the 
necessary  oxygen.     On  the  other  hand,  some  nitric  acid, 
placed  in  a  spoon  and  heated,  simply  boils  in  the  air,  but 
placed  in  the  hydrogen,  catches  fire  and  burns  with  a  pale 
yellow  flame.     The  nitric  acid  is  rich  in  oxygen,  and  the 
burning  is  the  combination  of  this  with  the  hydrogen  to 
form  water.     But  the  most  striking  results  are  obtained 
with  the  metallic  chlorates.    Potassium  chlorate  has  already 
been  referred  to  as  a  compound  of  potassium,  chlorine,  and 
oxygen.     The   other   chlorates    are   similarly   constituted. 
They  all  contain  a  metal,  chlorine,  and  much  oxygen,  and 
the  oxygen  is  only  loosely  held.     They  may  be  placed  in 
the  spoon,  heated  in  the  air,  and  plunged  into  the  hydrogen, 
where  they  burn  with  intense  light-display,  the  flame  hav- 
ing a  color  which  depends  on  the  particular  chlorate  em- 
ployed.    With    potassium    chlorate    it    is    violet,    barium 
chlorate  a  brilliant  green,  while  sodium  chlorate  gives  a 
yellow,  and  strontium  chlorate  a  red  flame   of  blinding 
brightness.     These  are  the  flame  colors  of  these  metals — 
that  is,  any  volatile  compound  of  the  metal,  introduced 
into  the  Bunsen  flame,  will  produce  the  corresponding  color, 
violet  for  potassium,  yellow  for  sodium,  and  so  on. 

63.  The  nature  of  flame. — From  these  experiments  we 
perceive  that  in  an  atmosphere  of  hydrogen  the  ordinary 


54  ELEMENTARY  CHEMISTRY 

phenomena  of  combustion  are  reversed :  substances  rich  in 
carbon  and  hydrogen  become  incombustible,  and  substances 
rich  in  loosely  held  oxygen  burn  brilliantly.  The  next  step 
is  to  inquire  into  the  nature  of  flame.  What  is  a  flame 
and  under  what  conditions  will  a  substance  burn  with 
flame?  We  have  already  met  with  cases  of  both  kinds  of 
combustion.  Thus,  sulphur,  phosphorus,  and  magnesium 
burn  in  oxygen  with  flame,  while  iron  and  charcoal  burn 
brilliantly,  but  without  flame.  A  flame  is  a  mass  of  highly 
heated  gas  or  vapor;  it  is  the  portion  of  space  in  which 
a  chemical  combination  is  occurring  with  great  evolution 
of  heat,  both  parties  to  the  combination  being  gases  or 
vapors  at  the  temperature  of  the  experiment. 

Iron  and  charcoal  burn  without  flame  because  they  are 
not  vaporized,  and  the  process  can  only  occur  at  the  surface 
of  the  iron  or  the  charcoal.  Magnesium  burns  with  flame, 
because  at  the  high  temperature  it  is  vaporized  and  the 
vapor  passes  into  the  space  immediately  surrounding  the 
burning  wire.  Thus,  the  combination  does  not  take  place 
at  the  surface  of  the  wire,  but  in  the  adjacent  space,  and 
there  is  produced  an  intensely  heated  mass  of  magnesium 
vapor,  oxygen,  and  magnesium  oxide,  which  make  up  the 
flame. 

Sulphur  and  phosphorus  behave  like  magnesium.  So 
with  a  candle,  the  wax  is  melted  and  drawn  up  by  the 
wick.  Approaching  near  the  heat,  it  is  converted  into  a 
mixture  of  combustible  gases. and  vapors;  and  these  burn 
with  a  flame.  It  will  be  clear  from  all  this  that  any  gas 
or  vapor  which  burns  at  all  will  burn  with  a  flame,  while  a 
solid  will  burn  with  flame  only  when  it  is  volatilized  at  the 
temperature  of  the  experiment. 

We  may  now  use  the  lamp-chimney  to  exemplify  an- 
other interesting  fact.  If  any  gas  A  will  burn  in  a  gas  B, 
then  B  will  also  burn  in  A  under  the  proper  conditions. 
A  tube  ending  in  a  fine  jet — the  ordinary  mouth  blow-pipe 


COMBUSTION 


55 


is  convenient — is  connected  with  a  cylinder  of  compressed 
oxygen,  and  a  very  gentle  stream  of  oxygen  allowed  to 
escape.  The  tube  is  now  intro- 
duced through  the  flame  of 
burning  hydrogen  into  the 
chimney.  The  oxygen  catches 
fire  from  the  hydrogen  flame 
and  goes  on  burning  in  the 
chimney  with  a  pale  blue  flame 
exactly  like  that  of  hydrogen 
burning  in  oxygen  (Fig.  20). 
It  is  easy  also  to  produce  a 
flame  of  air  burning  in  hydro- 
gen. In  order  to  do  this,  place 
in  the  bottom  of  the  chimney  a  FIG.  20.— combustion  of  oxygen  in 
rubber  cork  having  two  holes,  hydrogen, 

one  of  which  carries  the  hydrogen  tube,  a,  and  the  other  a 
short  rather  wide  glass  tube,  ~b,  open  at  both  ends.  With- 
draw the  cork  a  moment,  light 
the  hydrogen  at  its  tube,  and 
insert  the  cork  in  its  place.  The 
hydrogen  flame  soon  exhausts 
the  supply  of  oxygen  in  the 
vessel,  a  pale  blue  flame  floats 
about  the  chimney  for  an  in- 
stant, then  settles  on  the  open 
tube,  where  it  continues  to  burn. 
This  is  the  air-flame  (Fig.  21). 
We  may  now  light  the  hydro- 
gen above  at  the  hole  in  the  as- 
bestos plate,  and  we  have  above 
a  flame  of  hydrogen  burning  in 
air,  below  a  flame  of  air  burn- 
ing in  hydrogen.  The  latter  can  be  extinguished  instantly 
by  stopping  the  lower  end  of  the  air-tube  with  the  finger. 


FIG.  21.— Combustion  of  air  in 
hydrogen. 


56 


ELEMENTARY  CHEMISTRY 


64.  Structure  of  flame. — Into  the  flame  of  a  candle  we 
introduce  one  end  of  a  glass  tube  open  at  both  ends.    When 
the  tube  is  inclined  upward,  gases 
rise  through  it  and  may  be  lighted 

9 — f,n        m^\  H  at  the  end,  where  they  burn  with  a 

luminous  flame  (Fig.  22).  This  ex- 
periment— due  to  Faraday — shows 
that  the  candle  flame  is  hollow. 
This  is  true  of  all  flames.  The  hy- 
drogen flame  contains  an  inner  dark 
cone,  consisting  of  unburned  gas,  and 
an  outer  faintly  luminous  sheath 

which  is  the  p°rtion  of  sPace  in 

of  unbnrned  gases  ;    *,  lumi-  which    Combustion    is    taking    place, 
nous  portion ;  3,  outer  mantle,   m,  ,.  ,          i-t         r,    j 

when   complete  combustion  The    question    why    the   hydrogen 

takes  place.  flame    is    only    faintly    luminous, 

though  very  hot,  while  the  candle  or  gas  flame  gives  off  a 

bright  light,  has  been  much  discussed,  but  no  final  answer  has 

been  given.  It  is  pos- 
sible that  the  combus- 
tible gases  in  the  lat- 
ter case  are  decom- 
posed by  the  heat 


PIG.  23  A. — The  Bunsen  burner,    e  e, 
chimney ;  d,  holes  for  admission  of  air. 


Fie.  23  B.— Burner  with  chim- 
ney removed,  a  b,  base  ;  d, 
gas-supply  ;  c,  aperture  for 
admission  of  air. 


with  separation  of  solid  carbon  and  then  these  fine  solid  par- 
ticles, being  intensely  heated,  glow  brightly. 


COMBUSTION  57 

65.  The  Bunsen  burner. — The  coal-gas  flame  can  be 
made  non-luminous  by  depressing  a  flat-iron  or  other  cold 
metallic  object  upon  it.  In  the  same  way,  when  much  nitro- 
gen is  mixed  with  the  gas  before  it  is  burnt,  the  flame  be- 
comes blue  by  simple  cooling.  When  oxygen  is  mixed  with 
the  gas,  the  disappearance  of  the  light  is  due  to  a  different 
cause.  The  flame  becomes  much  hotter,  but  combustion  is 
so  perfect  that  no  carbon  separates  and  no  glowing  can  oc- 
cur. In  the  Bunsen  burner,  air  is  drawn  in  through  the 
holes  around  the  base,  a  mixture  of  gas  and  air  ascends  the 
tube  and  burns  at  the  top,  and  the  lack  of  luminosity  is 
probably  due  to  both  causes  acting  together :  to  the  dilution 
of  the  flame  by  the  nitrogen  and  the  greater  perfection  of 
combustion  brought  about  by  the  added  oxygen.  The  stu- 
dent will  best  understand  the  action  of  the  burner  by  taking 
it  apart  and  examining  it  with  some  care  (Fig.  23). 


CHAPTEK   VIII 

NAMING   CHEMICAL    COMPOUNDS— CHEMICAL   SYMBOLS  AND 
EQUATIONS— DIFFERENT  KINDS    OF   CHEMICAL    CHANGE 

66.  The  names  of  chemical  compounds. — Keturning  now 
to  the  combustions  in  oxygen  carried  out  in  the  sixth 
chapter,  we  have  seen  that  the  process  consists  in  the  rapid 
production  of  the  oxide  of  the  element  employed.  The 
termination  "  ide  "  generally  indicates  a  compound  of  two 
elements.  Thus,  an  oxide  consists  of  oxygen  in  combination 
with  some  other  element;  mercuric  oxide  consists  of  mer- 
cury and  oxygen.  In  the  same  way  a  sulphide  contains  sul- 
phur combined  with  something  else ;  lead  sulphide  is  a  com- 
pound of  lead  and  sulphur.  The  same  nomenclature  is  ap- 
plied to  other  elements.  Sodium  chloride,  common  salt,  is 
a  compound  of  sodium  and  chlorine ;  calcium  carbide,  used 
in  certain  bicycle  lanterns,  a  compound  of  calcium  and  car- 
bon, and  so  on. 

If  an  element  forms  more  than  one  oxide,  it  is  necessary 
to  distinguish  them  in  some  way  by  name.  Thus,  there 
are  two  compounds  of  mercury  with  oxygen.  One,  mer- 
curic oxide,  contains  200  parts  of  mercury  and  16  of  oxygen, 
the  other  400  parts  of  mercury  and  16  of  oxygen.  The  first 
we  have  already  learned  to  call  mercuric  oxide,  the  second 
is  black  and  is  called  mercurous  oxide.  There  are  two  com- 
pounds also  of  mercury  and  chlorine.  One,  the  well-known 
medicine  calomel,  contains  200  parts  of  mercury  and  35.5 
parts  of  chlorine.  This  is  mercurous  chloride.  The  other 
— corrosive  sublimate — contains  200  parts  of  mercury  and 
58 


NAMING  CHEMICAL  COMPOUNDS  59 

twice  35.5,  or  71  parts,  of  chlorine.  It  is  called  mercuric 
chloride. 

It  will  be  seen  from  this  that  the  ending  "  ous  "  is  ap- 
plied to  that  compound  which  contains  less  of  the  element 
coming  last  in  the  name  of  the  compound.  Mercurous 
bromide  contains  less  bromine  than  mercuric  bromide,  cu- 
prous sulphide  less  sulphur  than  cupric  sulphide. 

67.  Description  of  the  composition  of  compounds  by 
formulas. — Water  contains  2  parts  of  hydrogen  and  16  of 
oxygen  by  weight.  Let  us  agree  to  represent  1  part  of 
hydrogen  by  the  symbol  H.  Then  2  parts  can  be  repre- 
sented by  the  expression  H2.  Similarly,  16  parts  of  oxygen 
by  weight  can  be  represented  by  the  symbol  0.  Then  the 
composition  of  water  can  be  briefly  described  by  the  for- 
mula H20,  which  means  18  parts  of  water,  containing  16 
of  oxygen  and  2  of  hydrogen.  It  makes  no  difference  what 
particular  unit  of  weight  we  employ,  for  the  proportions 
are  relative  only,  but  it  will  fix  our  ideas  to  think  of  some 
definite  unit,  and,  since  the  gram  is  universally  employed 
in  science,  we  shall  use  that. 

There  are  453.6  grams  in  a  pound,  so  that  the  gram  is 
a  little  less  than  the  ^  of  an  ounce.  The  formula  H20, 
then,  means  18  grams  of  water,  of  which  2  grams  are  hydro- 
gen and  16  grams  oxygen.  Mercuric  oxide  we  have  seen  to 
contain  200  parts  of  mercury  and  16  of  oxygen.  Let  us 
represent  200  parts  of  mercury  by  the  symbol  Hg.  Then 
the  formula  HgO  is  a  concise  description  of  the  compo- 
sition of  the  substance.  It  signifies  216  grams  of  mercuric 
oxide,  of  which  200  grams  are  mercury  and  16  oxygen. 
Mercurous  oxide  contains  200x2,  or  400  grams  of  mercury 
and  16  of  oxygen  in  416  grams.  The  formula  becomes 
Hg,0.  So  the  symbol  Cl  is  employed  to  signify  35.5 
parts  by  weight — grams,  for  example — of  chlorine.  It  has 
just  been  mentioned  that  calomel,  mercurous  chloride,  con- 
tains 200  parts  of  mercury  and  35.5  of  chlorine.  Clearly 


60  ELEMENTARY  CHEMISTRY 

the  formula  is  HgCl.  Mercuric  chloride — corrosive  sub- 
limate— which  contains  35.5x2,  or  71  parts  of  chlorine  to 
200  of  mercury,  must  be  described  by  the  formula  HgCl2. 
Similarly,  the  symbol  of  sodium  is  Na,  and  the  weight 
which  experience  has  shown  to  be  most  suitable  for  de- 
scribing the  composition  of  sodium  compounds  is  23. 
NaCl  is  the  formula  of  table  salt,  and  this  means  that  in 
58.5  grams  of  it  there  are  23  grams  of  sodium  and  35.5 
of  chlorine. 

68.  Symbols  and  their  meaning. — Now  it  is  found  that 
by  means  of  the  symbol  H,  meaning  1  part  of  hydrogen  by 
weight,  the  composition  of  all  hydrogen  compounds  can  be 
briefly  described  in  formulas.     So  the  symbol  0,  meaning 
16  parts  of  oxygen,  enables  us  to  write  simple  formulas  like 
II20  for  water,  HgO  for  mercuric  oxide,  Hg20  for  mer- 
curous  oxide.    We  shall  find  as  we  progress  that  by  means 
of  this  symbol  the  composition  of  all  compounds  containing 
oxygen  can  be  described.     It  is,  we  may  say,  a  natural  chem- 
ical unit  in  which  oxygen  enters  into  its  compounds,  and 
while  we  often  have  to  multiply  it  in  order  to  express  -the 
composition  of  a  compound,  it  is  never  necessary  to  divide 
it.     In  the  same  way,  by  means  of  the  symbol  Hg,  mean- 
ing 200  parts  of  mercury,  the  composition  of  all  mercury 
compounds  without  exception  can  be  described.     Similarly 
with  Cl  =  35.5  parts  of  chlorine,  and  Na  =  23  parts  of 
sodium.     The  composition  of  every  known  compound  of 
these  elements  can  be  described  by  simple  formulas  based 
upon  these  quantities. 

69.  The  atomic  weights. — There  is  for  every  element  a 
natural  quantity  in  which  it  enters  into  its  compounds,  and 
by  means  of  which  concise  descriptions  of  the  composition 
of  these  compounds  can  be  given.    These  numbers  are  given 
in  the  table  in  the  appendix  and  are  called  the  "  atomic 
weights  "  of  the  elements.    This  term  is  derived  from  the 
hypothesis  that  material  things  are  composed  of  small  par- 


CHEMICAL  SYMBOLS  AND  EQUATIONS  61 

tides  with  vacant  spaces  between,  a  view  which  has  been 
helpful,  especially  in  dealing  with  the  very  numerous  com- 
pounds which  carbon  forms  with  other  elements.  It  will 
be  discussed  later.  At  present  it  is  important  for  the  stu- 
dent to  perceive  that  everything  which  has  been  stated  is 
pure  fact,  and  is  quite  independent  of  any  speculation  of 
this  kind. 

70.  Chemical  equations. — We  can  now  proceed  to  de- 
scribe chemical  changes  in  the  same  kind  of  language. 
When  by  heating  mercuric  oxide  it  separates  into  mercury 
and  oxygen,  we  write :  HgO  =  Hg  -J-  0,  which  means  that 
216  parts  of  mercuric  oxide  separate  into  200  of  mercury  and 
16  of  oxygen.  In  the  same  way  the  decomposition  of  water 
by  the  current  can  be  written :  H20  =  H2  +  0,  signifying 
that  18  grams  of  water  produce  2  of  hydrogen  and  16  of 
oxygen.  The  decomposition  of  potassium  chlorate  by  heat 
is  somewhat  more  complex.  Its  formula  is  KC103,  which, 
of  course,  means  39  parts  of  potassium,  35.5  of  chlorine 
and  16x3,  or  48  parts  of  oxygen,  in  combination.  When  it 
is  heated  the  change  is  represented  by  the  expression: 
KC103  =  KC1  +  30,  the  48  parts  of  oxygen  escaping,  while 
the  potassium  and  chlorine  remain  together,  as  74.5  parts 
of  potassium  chloride.  An  expression  of  this  kind  is  called 
a  chemical  equation.  It  is  simply  a  description  of  a  chemi- 
cal change,  in  a  kind  of  language  which,  for  clearness  and 
brevity,  has  probably  never  been  equaled.  The  sign  of 
equality  in  the  equation  stands  for  the  law  of  the  inde- 
structibility of  matter.  Since  nothing  is  lost  or  gained,  it 
stands  for  equality  in  weight. 

The  total  weight  of  the  substances  entering  into  the 
change — those  on  the  left-hand  side — must  be  equal  to  the 
total  weight  of  the  products — those  on  the  right-hand  side 
of  the  sign  of  equality.  Further,  no  element  has  ever  been 
transformed  into  any  other  element.  This  great  principle 
is  the  most  important  result  of  fifteen  hundred  years  of  un- 


62  ELEMENTARY  CHEMISTRY 

successful  efforts  by  the  alchemists  to  transform  copper, 
mercury,  lead,  and  other  cheaper  metals  into  gold  and  silver. 
Our  chemical  equations  must  conform  to  it.  Since  every 
symbol  represents  a  fixed  quantity  of  some  element  the 
number  of  symbols  corresponding  to  each  element  must  be 
the  same  on  both  sides.  Thus,  in  the  equation  for  the  de- 
composition of  potassium  chlorate,  we  have  on  each  side 
one  K,  one  Cl,  and  three  0.  This  equation  satisfies  our  two 
laws  therefore;  the  law  of  the  indestructibility  of  mat- 
ter and  that  which  expresses  the  impossibility  of  mutual 
transformation  of  the  elements,  and  any  equation  which 
violates  either  principle  is  absurd  by  inspection. 

An  equation  must  be  a  faithful  description  of  a  real 
process.  It  is  easy  to  write  equations  which  satisfy  both 
the  laws  mentioned  above  and  yet  are  false  because  the 
change  which  they  describe  does  not  really  happen.  The 
equation  KC103  =  KC10  +  02  is  a  case  of  this  kind.  So 
far  as  the  number  of  symbols  on  each  side  is  concerned, 
it  is  correct,  for  we  have  on  each  side  one  K,  one  Cl,  and 
three  0,  but  since  the  compound  KC10  is  not,  as  a  matter 
of  fact,  produced  when  potassium  chlorate  is  heated,  the 
equation  is  false.  In  the  same  way  every  correct  chemical 
formula,  like  HgO,  NaCl,  and  so  on,  represents  the  compo- 
sition of  a  real  chemical  compound.  Of  course,  the  formula 
can  not  be  written  until  the  compound  it  describes  has  been 
prepared  in  the  pure  state  and  analyzed.  It  is,  in  fact, 
one  way  of  writing  the  result  of  a  chemical  analysis  of  the 
compound. 

71.  Different  kinds  of  chemical  change. — When  magne- 
sium burns,  two  substances,  magnesium  and  oxygen,  dis- 
appear, and  one  substance,  magnesium  oxide,  is  produced. 
A  chemical  change  of  this  sort  is  called  a  combination. 
Other  examples  of  combination  are  the  union  of  copper 
and  sulphur  to  copper  sulphide  (p.  32)  and  of  hydrogen  and 
oxygen  to  form  water  (p.  7). 


DIFFERENT   KINDS  OF  CHEMICAL  CHANGE          63 

Combination  is  the  union  of  two  or  more  substances  to 
produce  one  substance. 

On  the  other  hand,  when  the  electric  current  is  caused 
to  pass  through  water,  the  water  separates  into  hydrogen 
and  oxygen.  This  is  the  reverse  of  combination  and  a 
change  of  this  kind  is  called  a  decomposition.  The  split- 
ting up  of  mercuric  oxide  into  mercury  and  oxygen  (p.  26) 
and  of  potassium  chlorate  into  potassium  chloride  and  oxy- 
gen (p.  41)  are  examples. 

Decomposition  is  the  separation  of  one  substance  into 
two  or  more. 

"  If  the  elements  should  cease  to  form  compounds  with 
each  other,  what  would  be  the  result.  All  chemical  com- 
pounds would  be  decomposed,  and  there  would  only  be  about 
eighty  different  kinds  of  substances.  All  living  things 
would  cease  to  exist,  arid  in  their  place  we  should  have 
three  invisible  gases  and  something  very  like  charcoal. 
Mountains  would  crumble  to  pieces,  and  all  water  would 
disappear,  giving  two  invisible  gases.  The  process  of  life 
in  its  many  forms  would  be  impossible."  (Eemsen.) 

There  are  other  important  kinds  of  chemical  change 
which  we  shall  become  familiar  with  as  we  meet  examples 
of  them  farther  on  in  our  work. 


CHAPTER   IX 


SALT  AND   SODIUM 

72.  Synthesis  of  salt. — A  fragment  of  sodium  is  placed 
in  a  bulb  blown  on  a  piece  of  hard  glass  tubing,  and  chlorine 
gas  is  passed  through  the  bulb.  On  being  gently  heated, 
the  sodium  catches  fire  in  the  chlorine  and  burns  with  a 
dazzling  yellow  flame.  When  the  bulb  has  cooled,  we  open 
it  and  find  a  white  powder  which  the  taste  shows  to  be 

common  salt  (Fig. 
24).  Thus  we  have 
shown,  by  synthe- 
sis, that  salt  is  a 
compound  of  sodi- 
um and  chlorine.1 
Its  formula  is 
NaCl. 

73.    Occurrence 
of    salt.  --  Sea- 
p».  ^.-synthesis  of  salt.  water.— Salt  is  the ' 

most  abundant  compound  of  sodium  as  well  as  the  most 
abundant  compound  of  chlorine,  and  it  is  important  be- 
cause it  is  the  raw  material  of  two  great  groups  of  chem- 
ical industries,  those  of  sodium  and  its  compounds,  and  those 
of  the  compounds  of  chlorine.  When  100  grams  of  the 
water  of  the  open  sea  are  evaporated  to  dryness,  3.5  grams 
of  solid  matter  are  left,  of  which  more  than  2.5  grams  are 

1  This  is  another  interesting  case  of  the  catalytic  action  of  water. 
Perfectly  dry  chlorine  does  not  act  visibly  upon  sodium. 
64 


SALT  AND  SODIUM  65 

sodium  chloride.  The  rest  consists  of  a  great  variety  of 
substances. 

Since  the  ocean  receives  the  washings  of  the  entire  crust 
of  the  earth,  and  since  everything  is  soluble  in  water  to  a 
greater  or  less  extent,  it  is  probable  that  sea-water  con- 
tains all  the  elements.  Nearly  half  of  them  have  already 
been  detected  in  it,  mostly  in  very  small  quantities.  The 
calcium  of  ocean-water  must  be  mentioned,  on  account  of 
its  importance  to  the  life  of  shell-fish;  magnesium  is  pres- 
ent in  considerable  quantity,  bromine  in  small  proportion, 
but  enough  to  make  its  extraction  profitable  under  certain 
conditions. 

Sea-water  does  not  appear  to  contain  any  dissolved 
iodine  compounds.  The  iodine,  which  it  always  contains,  is 
present  in  suspended  seaweed  spores. and  is  removed  by  fil- 
tering. The  larger  seaweeds  also  contain  iodine. 

Of  course,  sea-water  dissolves  gases  from  the  air,  oxygen, 
nitrogen,  and  carbon  dioxide,  and  these  are  important  to 
the  life  of  fishes  and  marine  plants.  Many  of  the  other  ele- 
ments are  present  in  traces  which  require  the  most  refined 
methods  of  chemical  analysis  to  detect  them.  Among  these 
we  may  mention  iron,  silver,  and  gold. 

74.  Water  of  inland  seas. — In  an  inland  sea  supplied  by 
streams  and  only  losing  water  by  evaporation,  the  quan- 
tity of  dissolved  matter  may  become  very  great,  for  the 
water  which  enters  contains  dissolved  solids,  and  that  which 
leaves  contains  none.     This  is  the  state  of  the  Great  Salt 
Lake  and  the  Dead  Sea.     One  hundred  parts  of  the  water 
of  the  latter,  when  evaporated,  leave  nearly  23  parts  of  solid 
residue,  consisting  chiefly  of  the  chlorides  of  magnesium, 
calcium,  and  sodium,  and  there  are  other  lakes  whose  waters 
contain  as  much  as  30  per  cent  of  material  in  solution.     In 
all  water  of  this  kind  marine  life  is  impossible. 

75.  Commercial  production  of  salt. — In  Siberia  some  salt 
is  obtained  from  sea-water  by  freezing.     The  ice  which  sep- 


66  ELEMENTARY  CHEMISTRY 

arates  is  fresh.  It  is  removed,  and  the  freezing  continued 
until  the  solution  of  salt  becomes  strong  enough  to  boil 
down  profitably.  Salt  is  obtained  in  warm  climates,  espe- 
cially on  the  shores  of  the  Mediterranean,  by  natural  evap- 
oration. Sea-water  is  allowed  to  run  into  shallow  basins 
in  which  it  evaporates  of  itself  until  the  salt  crystallizes 
out.  Kock-salt  beds  have  probably  been  produced  by  the 
cutting  off  of  arms  of  the  sea  from  the  main  body,  during 
the  geological  past.  Evaporation  would  then  occur,  and 
the  end  result  would  be  a  bed  of  salt  which  might  become 
covered  with  a  water-tight  layer  of  clay,  and  so  be  protected 
from  the  dissolving  action  of  the  rains. 

England  produces  more  salt  from  rock-salt  beds  than 
any  other  country,  and  there  are  deposits  in  Austria  and 
Germany  and  in  various  parts  of  the  United  States,  for  in- 
stance in  southwestern  New  York. 

76.  Salt    springs. — In    the    same    districts    occur    salt 
springs.     The   evaporation   of  such  natural   brine,   or   of 
artificial  brine  made  by  turning  water  into  a  boring  in  a 
salt  bed,  is  the  most  important  source  of  salt.     The  water 
must  contain  at  least  16  per  cent  of  salt  in  order  to  be 
evaporated  profitably.    It  is  boiled  down  in  iron  pans  until 
the  salt  separates. 

77.  Properties  of  salt. — Eock  salt  is  often  colored  blue 
or  yellow,  but  pure  sodium  chloride  is  colorless.    It  crystal- 
lizes in  cubes.     It  dissolves  in  about  twice  its  weight  of 
water,  and  the  solubility  increases  only  very  slightly  when 
the  liquid   is  heated.     It  is   almost   insoluble  in   alcohol. 
At  a  red  heat  it  melts  and  at  a  slightly  higher  temperature 
rapidly  vaporizes.     This  can  easily  be  seen  by  sprinkling 
some  salt  on  a  hot  fire. 

78.  Function  of  salt  in  the  diet. — Salt  is  always  used  by 
civilized  man  at  the  table,  and  there  has  been  some  question 
as  to  whether  it  is  a  mere  flavoring  material,  like  pepper, 
or  is  a  necessary  element  in  a  healthful  diet.     The  facts 


SALT  AND  SODIUM  67 

seem  to  favor  the  latter  statement.  The  gastric  juice 
always  contains  a  little  hydrochloric  acid,  HC1,  and  this 
is  formed  from  the  salt  of  the  food.  Experiments  on 
dogs  have  shown  that,  when  the  body  is  deprived  of  salt 
for  a  long  time,  this  hydrochloric  acid  disappears  and  the 
functions  of  the  stomach  are  no  longer  rightly  performed. 
Other  chlorides  can  serve  the  same  purpose  as  sodium 
chloride  in  the  diet.  There  are  African  tribes  which  em- 
ploy potassium  chloride,  KC1,  in  a  similar  way.  Salt,  in 
enormous  quantities,  acts  as  an  irritant  poison.  In  one 
case  250  grams  (about  one-half  pound),  taken  at  one  time, 
caused  death. 

79.  If  a  dilute  solution  of  sodium  chloride  is  cooled  in  a  freezing 
mixture  and  stirred  with  a  thermometer,  it  is  found  to  begin  to 
freeze  at  a  lower  temperature  than  pure  water,  and  the  ice  which 
separates  is  fresh.  Now  let  the  cooling  be  continued.  For  a  while 
ice  continues  to  separate.  Finally,  there  comes  a  time  when  the 
remaining  liquid  freezes  as  a  whole,  ice  and  salt  separating  together. 
From  the  moment  the  first  crystal  of  salt  separates  along  with  the 
ice  the  temperature  remains  constant  until  the  whole  mass  has  be- 
come solid. 

Most  solutions  behave  in  this  way.  When  they  are  cooled,  pure 
ice  separates  first,  and  the  temperature  falls  until  the  separation  of 
solid  dissolved  substance  begins.  Then  the  temperature  remains  the 
same  until  the  solidification  is  complete.  Solutions  in  liquids  other 
than  water  act  similarly. 

In  the  same  way,  if  a  salt  solution  is  heated,  it  begins  to  boil  at  a 
higher  temperature  than  pure  water.  The  vapor  of  water  is  given  off 
and  this  leaves  behind  it  a  stronger  solution.  Now,  the  stronger  the 
solution  the  higher  the  boiling  point,  so  that  the  temperature  rises. 
Finally,  the  continued  loss  of  steam  leaves  the  solution  saturated, 
and  solid  salt  separates.  From  this  moment  the  temperature  is 
fixed  until  the  solution  has  been  evaporated  to  dryness.  This  leaves 
solid  salt,  which,  of  course,  can  be  heated  to  any  desired  degree, 
This  behavior  also  is  common  to  most  solutions. 


68  ELEMENTARY  CHEMISTRY 

SODIUM 

80.  Historical. — In  discussing  hydrogen,  it  has  been  re- 
marked that  when  sodium  acts  upon  water  it  liberates  half 
of  the  hydrogen  and  combines  with  the  other  half  and  with 
all  the  oxygen,  forming  a  compound  called  sodium  hydrox- 
ide.   We  can  now  describe  this  process  in  symbols. 

Na  +  H20  =  NaOH  +  H. 

The  substance  which  has  the  composition  NaOH — sodium 
hydroxide — is  familiarly  called  concentrated  lye  or  caustic 
soda.  It  was  for  many  years  supposed  to  be  an  element, 
but  in  1807  Sir  Humphry  Davy  announced  that  when  the 
electric  current  is  caused  to  pass  through  a  mass  of  mois- 
tened caustic  soda,  small  globules  of  metal  collect  at  the 
negative  pole.  From  this  he  concluded  that  caustic  soda 
was  a  compound  containing  a  metal,  and  proposed  for  this 
metal  the  name  sodium. 

81.  Preparation  of  sodium. — After  trying  many  other 
methods,  chemists  have  finally  returned  to  a  modification 
of  Davy's  process  as  the  best  for  making  sodium  on  a  com- 
mercial  scale.     The   current    from    a    dynamo   is    passed 
through  caustic  soda,  heated  just  to  melting  in  an  iron 
vessel.    Oxygen  is  liberated  at  the  positive  pole  and  hydro- 
gen and  sodium  together  at  the  negative  pole : 

NaOH  =  Na  +  0  +  H. 

Many  attempts  have  been  made  to  prepare  sodium  by 
the  action  of  the  current  upon  salt,  which  is  far  cheaper 
than  caustic  soda.  So  far  the  methods  have  not  succeeded 
in  practice,  but  the  efforts  still  continue  and  may  at  any 
time  be  successful. 

82.  Occurrence  of  sodium  compounds. — Sodium  chloride 
is  the  most  common  compound  of  the  metal,  but  other  sodi- 
um compounds  are  abundant  in  nature.     Sodium  nitrate, 


SIR  HUMPHRY  DAVY 
B.  England,  1778.     D.  1829. 


SALT  AND  SODIUM  69 

NaN03,  occurs  in  great  deposits  on  the  western  coast  of 
South  America.  Most  of  the  common  rocks  contain  sodium 
compounds.  It  stands  seventh  among  the  elements  in  point 
of  abundance,  and  recent  calculations  show  that  the  acces- 
sible portion  of  the  earth's  crust  contains  about  2J  per 
cent  of  it. 

83.  Properties. — Sodium  is   a   metal   with  a  brilliant, 
somewhat  pink  luster,  which  it  instantly  loses  in  the  air. 
This  is  due  to  the  fact  that  the  water-vapor  in  the  air  acts 
upon  it  in  just  the  same  way  that  liquid  water  does,  con- 
verting the  surface  into  sodium  hydroxide.    Sealed  up  in  a 
glass  tube,  from  which  the  air  has  been  removed  by  a  cur- 
rent of  dry  hydrogen,  the  luster  is  permanent.     It  is  soft 
like  wax,  can  be  cut  readily  with  a  knife,  and  is  an  excel- 
lent conductor  of  electricity  and  heat. 

Heated  in  the  absence  of  air,  it  melts  to  a  mercury-like 
liquid,  which,  when  further  heated,  passes  into  a  purple 
vapor.  Heated  in  air  or  oxygen,  it  burns  with  a  yellow 
flame. 

84.  Oxides  of  sodium. — Two  oxides  can  be  produced  ac- 
cording to  the  circumstances.     If  the  temperature  is  high 
and  the  supply  of  oxygen  small,  the  product  is  sodium  mon- 
oxide, a  gray  substance,  of  the  composition  Na20.     When 
this  substance  is  sprinkled  with  water,  it  combines  violently 
with  it,  producing  sodium  hydroxide: 

Na20  +  H20  =  2NaOH. 

If,  on  the  other  hand,  sodium  is  heated  to  a  low  tempera- 
ture with  abundant  air  supply,  the  product  is  yellow  sodium 
peroxide,  Na202.  This  is  made  on  a  large  scale  in  this  way, 
and  is  used  for  bleaching,  because,  when  treated  with  a 
dilute  acid,  it  yields  a  solution  containing  hydrogen  per- 
oxide. Hydrogen  peroxide  has  the  formula  H202,  and  it  dif- 
fers from  water,  therefore,  in  containing  32  parts  of  oxygen 
instead  of  16,  in  combination  with  2  parts  of  hydrogen. 


70  ELEMENTARY  CHEMISTRY 

Its  production  from  sodium  peroxide  and  hydrochloric  acid 
is  represented  thus  : 

Na202  +  2HC1  =  2NaCl  +  HA-1 

85.  Sodium  hydroxide,  NaOH,  can  be  obtained  pure  by 
treating  sodium  with  water  in  a  silver  dish: 

Na  +  H20  =  NaOH  +  H. 

Since  the  action  is  violent,  only  a  small  fragment  of  the 
metal  is  brought  into  contact  with  the  water  at  once.  The 
liquid  is  finally  evaporated  to  dryness  and  the  residue 
melted.  This  method  is  employed  in  the  laboratory,  but 
it  is  far  too  expensive  for  the  preparation  of  the  vast  quan- 
tities of  sodium  hydroxide  which  are  made  industrially. 
Practically,  it  is  made  from  washing  soda,  sodium  carbon- 
ate, Na2C03.  This  is  dissolved  in  water  and  the  liquid 
boiled  with  slaked  lime,  calcium  hydroxide  Ca  (OH)2,  in 
iron  vessels.  The  reaction  which  takes  place  is  expressed 
by  the  equation 

Na2C03  +  Ca(OH)2  =  CaC03 


Calcium  carbonate,  CaC03,  is  the  same  thing,  chemically, 
as  marble  or  chalk.  It  is  insoluble  in  water  and  settles  to 
the  bottom,  so  that  the  clear  solution  of  sodium  hydroxide 
can  be  poured  off  from  it.  This  is  then  evaporated  to  dry- 
ness,  melted,  and  cast  into  sticks.  Chemists  are  actively 
at  work  on  the  problem  of  producing  sodium  hydroxide 
directly  from  salt  by  the  electrical  method,  and  it  is  prob- 
able that  this  method  will  be  perfected  and  will  displace 
the  other  in  the  near  future. 

86.  Properties.  —  Sodium  hydroxide  is  usually  sold  in  the 
form  of  sticks.     It  is  a  tough,  white,  translucent  mass, 

1  Of  course  the  formulae  HO  for  hydrogen  peroxide  and  NaO  for 
Isodium  peroxide  would  represent  exactly  the  same  composition  by 
weight.  The  reasons  for  using  doubled  formulas  will  be  given  later. 


SALT  AND  SODIUM  71 

which  can  be  melted  and  vaporized  without  decomposition. 
It  is  very  soluble  in  water,  and  the  solution  has  a  slippery 
feel,  because  it  acts  chemically  upon  the  oil  of  the  skin, 
producing  a  soap. 

87.  Deliquescence. — When  a  stick  of  sodium  hydroxide 
is  allowed  to  remain  exposed  to  the  air,  it  becomes  wet,  be- 
cause water-vapor  is  attracted  from  the  air  and  the  solution 
of  sodium  hydroxide  is  formed.    Many  other  substances  be- 
have similarly,  and  this  occurrence  is  called  deliquescence. 
Whether  a  substance  attracts  water  from  the  air  and  dis- 
solves in  this  water  depends  entirely  upon  how  much  water 
the  air  contains  at  the  time.     Any  soluble  substance  will 
deliquesce  when  the  air  is  saturated  with  water-vapor.    It 
is  a  familiar  fact  that  table  salt,  which  is  not  deliquescent 
ordinarily,  becomes  deliquescent  at  the  seashore  on  account 
of  the  large  quantity  of  moisture  in  the  air. 

88.  Uses  of  sodium  hydroxide. — Sodium  hydroxide  is  a 
commercial  product  of  great  importance.    It  is  used  in  enor- 
mous quantities  for  the  production  of  soap,  which  is  made 
by  boiling  caustic-soda  solution  with  some  oil  or  fat.     In 
the  case  of  castile  soap,  olive  oil  is  used,  palm  oil  for  palm 
soap,  and,  for  cheaper  soaps,  animal  fats.     Sodium  hydrox- 
ide and  its  solution  absorb  carbon  dioxide  gas,  C02,  greedily 
from  gaseous  mixtures  containing  it,  and  are  employed  in 
the  laboratory  for  this  purpose.    In  this  reaction  the  sodium 
hydroxide  passes  into  sodium  carbonate, 

SNaOH  +  C02  =  Na2CO,  4-  H20. 

89.  Sodium  carbonate,  N"a2C03,  called  soda  or  washing 
soda,  is  made  from  common  salt  by  methods  which  can  not 
be  fully  discussed  here.     One  process  consists  in  dissolving 
the  sodium  chloride  in  water  and  passing  into  the  solution 
first  ammonia  gas,  NH3,  and  then   carbon  dioxide,   CO... 
This  is  the  equation, 

NaCl  +  C02  +  NH3  +  H20  =  NaHC03  +  NH4C1. 


72  ELEMENTARY  CHEMISTRY 

The  substance  NaHC03,  called  mono-sodium  carbonate  be- 
cause it  contains  only  one  atomic  weight  of  sodium,  is 
ordinary  baking-soda.  Not  being  very  soluble  in  water,  it 
separates  as  a  white  powder,  which  can  be  converted  into 
sodium  carbonate,  Na2C03,  by  heat, 

2NaHC03  =  Na2C03  +  H20  +  C02. 

The  carbon  dioxide  liberated  in  this  change  is  used  again 
in  treating  more  sodium  chloride.  The  substance  having  the 
composition  NH4C1,  which  is  the  other  product  of  the  treat- 
ment of  the  sodium  chloride,  is  called  sal-ammoniac  or  am- 
monium chloride.  We  shall  study  it  later,  but  it  is  neces- 
sary to  remark  now  that  the  ammonia  can  be  obtained  from 
it  again,  and  used  over  and  over. 

Sodium  carbonate  is  a  white  powder,  which  melts  at  a 
red  heat,  and  vaporizes  undecomposed  at  a  higher  tempera- 
ture. It  has  a  bitter,  nauseous  taste,  and  .is  soluble  in  about 
five  times  its  weight  of  water  at  ordinary  temperatures. 
It  is  much  more  soluble  in  hot  water,  and  when  a  hot, 
strong  solution  is  cooled,  there  separates  not  the  white  pow- 
der which  was  originally  dissolved,  but  colorless  crystals, 
which  can  be  obtained  quite  large  by  slow  cooling.  If  one 
of  the  crystals  is  dried  carefully  with  blotting-paper  and 
heated  in  a  dry  test-tube,  it  will  again  be  converted  into  the 
white  powder,  and  much  water  will  condense  in  the  cold 
upper  part  of  the  tube.  By  making  a  similar  experiment 
with  a  weighed  quantity  in  a  platinum  dish,  it  can  be  shown 
that  the  quantity  of  sodium  carbonate  indicated  by  the  for- 
mula Na2C03  (106  parts)  is  in  combination  in  the  crystals 
with  180  parts  of  water,  ten  times  the  quantity  represented 
by  the  formula  H20.  Hence  we  write  the  formula  of  crys- 
tallized sodium  carbonate  Na2C0310H20.  This  is  the  wash- 
ing-soda of  the  household. 

90.  Water  of  crystallization. — Water  in  this  condition, 
chemically  combined  with  a  salt,  is  called  water  of  crystal- 


SALT  AND  SODIUM  73 

lization.  When  it  is  driven  out  the  crystal  falls  to  pieces. 
The  same  salt  may,  and  usually  does,  form  several  com- 
pounds with  water  in  different  proportions.  Several  are 
known  in  the  case  of  sodium  carbonate,  of  which  one, 
Na2C03.H20,  is  becoming  an  important  technical  product. 
Many  other  liquids  play  a  similar  role.  Thus  we  have  alco- 
hol of  crystallization,  chloroform  of  crystallization,  and  so 
on.  Some  substances,  salt  and  potassium  chlorate,  for  in- 
stance, separate  from  aqueous  solution  in  crystals  which 
contain  no  water  and  are  said  to  be  anhydrous. 

When  crystallized  sodium  carbonate  is  exposed  to  air  it 
falls  slowly  to  a  white  powder,  which  has  the  composition 
Na2C03.H20,  nine-tenths  of  the  water  being  given  off — 

Na2C0310H20  =  Na2C03.H20  +  9H20. 

91.  Efflorescence. — A  substance  which  loses  its  water  of 
crystallization  in  the  air,  falling  to  a  powder  in  this  way, 
is  said  to  effloresce.  This  phenomenon,  like  deliquescence, 
depends  on  the  quantity  of  water-vapor  in  the  air.  If  trans- 
ported to  the  Sahara,  many  crystallized  salts  which  we  call 
permanent  would  at  once  be  classed  as  efflorescent.  And  in 
air  saturated  with  water-vapor,  Na2C0310H20,  will  not 
effloresce.  It  will  deliquesce,  absorbing  water-vapor  and 
producing  a  solution. 

Mono-sodium  carbonate,  NaHC03,  baking-soda,  is  a 
white  crystalline  powder.  Its  decomposition  by  heat  into 
Na2C03,  water  and  carbon  dioxide,  has  already  been  referred 
to.  It  is  not  used  alone  in  the  baking  process  because  the 
sodium  carbonate  which  would  be  left  in  the  bread  would 
make  it  nauseous  and  unwholesome.  We  place  a  little 
baking-soda  in  a  beaker  and  pour  some  hydrochloric  acid 
over  it.  A  violent  escape  of  carbon  dioxide  takes  place : 
NaHCO.,  +  HC1  =  NaCl  +  H20  +  C02. 

Thus,  when  an  acid  acts  upon  mono-sodium  carbonate, 
the  sodium  salt  of  this  acid — that  is,  a  compound  in  which 


74  ELEMENTARY  CHEMISTRY 

the  hydrogen  which  acids  always  contain  is  replaced  by 
sodium — is  produced,  together  with  water,  and  carbon  diox- 
ide escapes. 

92.  Baking-powder. — A  baking-powder  is  always  a  mix- 
ture of  mono-sodium  carbonate  with  some  substance  which 
acts  upon  it  like  an  acid,  liberating  carbon  dioxide  from  it. 
In  the  best  powders,  cream  of  tartar  is  employed  for  this 
purpose,  while  in  cheaper  powders  alum  is  used. 

93.  Sodium  sulphate,  Na2S04,  is  made  by  the  action  of 
sulphuric  acid  on  salt — 

SNaCl  +  H2S04  =  Na2S04  +  2HC1. 

It  separates  from  solution  in  water  in  large,  colorless  crys- 
tals of  the  composition  Na2S0410H20,  which  effloresce  in 
dry  air.  These  crystals  were  formerly  supposed  to  have 
wonderful  medicinal  qualities. 


CHAPTER   X 

CHLORINE 

94.  Preparation. — Some  coarsely  powdered  manganese 
dioxide  is  placed  in  a  flask  provided  with  a  rubber  cork  in 
which  are  two  perforations.  Through  one  hole  passes  a 
tube  having  a  funnel  at  the  upper  end ;  the  lower  end  runs 
nearly  to  the  bottom  of  the  flask.  The  other  hole  carries 
a  tube  which  runs  just  to  the  inside  of  the  cork.  This 
tube — called  the  delivery  tube — is  bent  twice  at  right  angles, 
and  runs  to  the  bottom  of  a  large,  empty  jar.  Chlorine  gas, 
being  about  two  and  a  half  times  as  heavy  as  air,  flows  into 
the  bottom  of  the  jar  like  water  and  displaces  the  air,  and 
this  is  the  best  way  to  collect  it,  for  it  is  soluble  in  water 
and  can  not  be  collected  over  it. 

Strong  hydrochloric  acid  is  poured  through  the  funnel 
tube  until  about  six  times  as  much  by  weight  has  been  in- 
troduced as  of  manganese  dioxide.  Then  a  small  flame  is 
placed  under  the  flask,  which  has  been  supported  on  wire 
gauze  to  avoid  cracking.  The  greenish-yellow  color  of 
chlorine  rapidly  appears  in  the  generating  flask,  and  then  in 
the  bottom  of  the  collecting  jar,  from  which  it  rises  to  the 
top.  Then  the  tube  is  withdrawn,  the  ground-glass  cover, 
which  has  been  resting  loosely  on  the  jar,  is  tightly  placed 
over  it,  the  jar,  which  is  now  full  of  chlorine,  is  stood  aside, 
and  another  takes  its  place.  In  this  way  a  number  of  jars 
full  of  the  gas  can  be  collected,  but  great  care  must  be  taken 
not  to  inhale  the  gas,  as  it  is  very  injurious.  The  appara- 
tus employed  is  shown  in  Fig.  25. 

75 


76 


ELEMENTARY  CHEMISTRY 


Chlorine  was  first  obtained  in  1774  by  the  great  Swedish 
chemist,  Carl  Wilhelm  Scheele,  who  prepared  it  by  the 

method  just  de- 
scribed. There  are 
dozens  of  ways  in 
which  it  can  be 
made  in  the  labora- 
tory, but  it  hap- 
pens that  this  same 
process  is  employed 
industrially  for  the 
production  of  the 
gas  on  a  large 
scale.1  Instead  of 
glass  flasks,  large 
vessels  of  stone- 
ware, or  of  flag- 
stones bolted  to- 
gether, are  used. 
The  chemical  change  which  takes  place  in  our  flask  is  de- 
scribed by  the  equation : 

Mn02  +  4HC1  =  MnCl2  +  2H20  +  C12. 

95.  Physical  properties. — Chlorine  is  a  greenish-yellow 
gas,  which  has  been  converted  by  cold  and  pressure  into  a 
clear  yellow  liquid  heavier  than  water.  The  gas  is  nearly 
35.5  times  as  heavy  as  hydrogen,  and  therefore  about  2J 
times  as  heavy  as  air.  It  has  a  suffocating  smell,  and  its 
inhalation  gives  rise  to  dangerous  inflammation  of  the  mu- 
cous membranes  of  the  respiratory  passages,  and  may  cause 
death.  Even  chlorine  very  much  diluted  with  air  is  injuri- 

1  At  present  this  method  of  making  chlorine  on  the  large  scale  is 
being  replaced  by  the  electrolysis  of  solutions  of  potassium  chloride, 
KC1,  or  of  sodium  chloride,  NaCl.  When  the  electric  current  is  passed 
through  such  a  liquid,  chlorine  is  liberated  at  the  positive  pole. 


FIG.  25.— Preparation  of  chlorine  from  manganese 
dioxide  and  hydrochloric  acid. 


CHLORINE  77 

ous,  though  experiments  on  dogs  have  shown  that  the  quan- 
tity which  can  be  borne  without  injury  increases  with  habit. 
Water  dissolves  about  twice  its  volume  of  the  gas  at  ordi- 
nary temperatures.  This  solution  is  called  chlorine"  water, 
and  when  it  is  cooled  there  separates  from  it  in  yellow 
scales  a  compound  of  chlorine  with  water  of  crystallization 
called  chlorine  hydrate,  C12  8H20.  This  separates  again 
into  chlorine  and  water,  slowly,  if  kept  in  a  freezing  mix- 
ture, and  rapidly  at  room  temperature. 

96.  Chemical   properties. — Chemically,   chlorine   is   ex- 
tremely energetic,  easily  entering  into  combination  with 
many  other  elements;  this  is  the  reason  it  is  not  found  as 
such  in  nature.     When  powdered  arsenic  is  sprinkled  into 
the  gas  it  catches  fire  and  burns  to  arsenic  chloride,  AsCl3, 
a  colorless  poisonous  liquid  which  floats  in  the  jar  for  a 
time  as  a  smoke.    Powdered  antimony  behaves  in  the  same 
way.     A  piece  of  phosphorus  introduced  into  the  gas  in  a 
spoon  melts  and  then  takes  fire,  burning  with  a  pale  flame 
to  phosphorus  trichloride,  PC13,  which  then,  if  there  is 
plenty  of  chlorine,  combines  with  more  of  it,  producing 
phosphorus  pentachloride,  PC15.     Chlorine  has  no  action 
upon  carbon,  and  if  a  piece  of  glowing  charcoal  is  plunged 
into  the  gas  it  is  extinguished.1     If  a  burning  candle  is 
plunged  into  chlorine  the  flame  becomes  red  and  feeble,  and 
quantities  of  soot  separate  on  the  glass.     The  hydrogen  of 
the  candle  burns  in  the  chlorine  to  hydrochloric  acid,  HC1, 
but  the  carbon  separates  in  the  free  state. 

97.  Action  of  chlorine  upon  the  metals. — Chlorine  acts 
upon  the  metals,  converting  them  into  chlorides,  in  many 
cases  with  evolution  of  light  and  heat.    The  brilliant  com- 
bustion of  sodium  has  already  been  described.    Copper  leaf 

1  It  must  not  be  concluded  from  this  that  no  compounds  of  chlorine 
with  carbon  are  known.  In  most  cases  where  two  elements  do  not 
combine  directly  their  compounds  can  be  obtained  by  indirect  meth- 
ods, and  it  is  so  in  the  present  instance. 


78  ELEMENTARY  CHEMISTRY 

— Dutch  leaf,  as  it  is  called — takes  fire,  producing  copper 
chloride.  A  brass  wire,  whose  combustion  can  be  started 
by  placing  a  little  Dutch  leaf  on  the  end,  burns  energetically, 
producing  the  chlorides  of  the  metals  of  which  brass  con- 
sists— copper  and  zinc. 

Some  of  these  combustions  are  not  interfered  with  by 
drying  the  chlorine,  those  of  arsenic  and  antimony,  for  ex- 
ample. In  other  cases  the  action  does  not  appear  to  occur 
at  all  when  the  gas  is  completely  dry.  Sodium  and  copper 
retain  their  luster  in  dry  chlorine  for  years.  In  these  cases, 
and  in  all  other  similar  ones,  it  is  probable  that  the  pres- 
ence of  the  water  simply  increases  greatly  the  speed  of  a 
process  which  would  occur,  and  yield  the  same  products,  in 
its  absence.  This,  in  fact,  is  what  we  mean  by  catalytic 
action.  Copper,  for  example,  combines  with  dry  chlorine, 
but  the  change  is  so  slow  that  it  escapes  detection:  On 
addition  of  water-vapor  the  combination  becomes  very  rapid, 
heat  is  evolved  rapidly,  the  temperature  rises  high  enough 
to  make  the  products  luminous,  and  we  call  the  process  a 
combustion. 

98.  Combustion  of  hydrogen  in  chlorine. — A  hydrogen 
flame,  fed  by  a  generator  containing  zinc  and  dilute  sul- 
phuric acid,  is  lowered  into  a  jar  of  chlorine.  It  continues 
to  burn,  but  becomes  large  and  pale.  After  a  time  the  flame 
goes  out,  and  then  it  is  found  that  the  chlorine  has  disap- 
peared and  the  jar  is  filled  with  a  colorless  gas  of  quite 
different  properties.  This  is  hydrochloric  acid,  HC1.  When 
equal  volumes  of  hydrogen  and  chlorine  are  mixed  in  the 
dark  and  exposed  to  sunlight  or  to  the  light  of  burning 
magnesium,  explosion  takes  place  and  hydrochloric  acid  is 
formed.  If  both  gases  are  completely  free  from  water  no 
explosion  occurs,  nor  does  the  mixture  explode  if  kept  very 
cold — at  — 12°,  for  instance.  If  the  mixture  is  exposed 
to  ordinary  diffuse  daylight,  slow  combination  without  ex- 
plosion takes  place. 


CHLORINE  79 

99.  Action   of   chlorine   upon  hydrogen   compounds.— 
Bleaching. — Chlorine   even   attacks   hydrogen   compounds, 
removing  some  or  all  of  the  hydrogen  to  form  hydrochloric 
acid.    In  this  way  it  slowly  decomposes  water  in  sunlight : 

H20  +  C12  =  2HC1  +  0. 

Upon  these  facts  depends  the  bleaching  action  of  chlo- 
rine, for  all  vegetable  and  animal,  and  many  artificial  color- 
ing matters,  are  compounds  of  hydrogen.  Either  the  chlo- 
rine attacks  the  coloring  matter  and  destroys  it  by  remov- 
ing the  hydrogen  to  produce  hydrochloric  acid,  or  else  the 
chlorine  combines  with  the  hydrogen  of  the  water  which  is 
always  present  in.  bleaching  operations,  and  the  oxygen  lib- 
erated destroys  the  coloring  matter.  Perfectly  dry  chlorine 
does  not  bleach. 

100.  Hydrochloric   acid,   HC1. — A   small,   strong  glass 
tube,  narrow  at  both  ends  and  provided  at  each  end  with  a 
stop-cock,  is  covered  with  a  wire-gauze  jacket,  to  avoid  dan- 
ger to  the  eyes  in  case  of  breakage,  and  filled  in  the  dark 
with  a  mixture  of  equal  parts  of  hydrogen  and  chlorine. 
The  tube  employed  is  that  shown  in  Fig.  29.     The  other 
apparatus  in  the  figure  is  not  needed  in  this  experiment. 
The  mixture  is  caused  to  explode  by  burning  magnesium 
wire  near  it.    The  tube  now  contains  hydrochloric  acid.    It 
is  allowed  to  cool  perfectly,  one  end  is  placed  under  the 
surface  of  some  mercury  in  a  dish,  and  the  stop-cock  at  that 
end  opened.     No  gas  escapes  and  no  mercury  enters.    This 
shows  that  the  volume  of  the  hydrochloric  acid  produced 
is  equal  to  the  sum  of  the  volumes  of  the  hydrogen  and  the 
chlorine,  or — 

1  volume  hydrogen  +  1  volume  chlorine  —  2  volumes 
hydrochloric  acid. 

101.  Combination  of  gases  by  volume. — It  is  important 
to  notice  the  simplicity  of  these  relations.     It  will  be  re- 


80  ELEMENTARY  CHEMISTRY 

called  that  we  found  a  similar  state  of  things  in  the  com- 
bination of  hydrogen  and  oxygen  to  steam: 

2  volumes  hydrogen  +  1  volume  oxygen  =  2  volumes  steam. 

This  is  always  the  case  when  combination  occurs  be- 
tween gases.  There  is  always  a  simple  relation  between 
the  volumes  of  the  gases  which  combine.  And  if  the  com- 
pound is  also  a  gas,  there  is  also  a  simple  relation  be- 
tween the  volume  of  each  gas  and  that  of  the  compound. 
Precisely  the  same  thing  is  true  of  decompositions  in 
which  gases  are  produced.  This  fact  is  usually  called  the 
law  of  simple  volume  ratios.  It  is  true  also  of  vapors. 

102.  Electrolysis  of  hydrochloric  acid. — Our  conclusion, 
that  hydrochloric  acid  contains  equal  volumes  of  hydrogen 
and  chlorine,  can  be  corroborated  by  decomposing  its  aque- 
ous solution  by  the  electric  current.     For  this   purpose, 
the    same   apparatus    can    be    employed    which    served    in 
the   electrolysis   of  water    (Fig.    2).     When   the   current 
passes,  hydrogen  separates  at  the  negative  pole  and  chlo- 
rine at  the  positive.     It  is  necessary  to  allow  the  current 
to  pass  half  an  hour  before  measuring  the  gases.     This 
is  on  account  of  the  solubility  of  chlorine.     Some  of  it 
dissolves  at  first,  and  the  volume  of  the  gas  collected  is 
smaller  than   that   of  the   hydrogen,   owing   to   this   loss. 
Finally,   the   liquid   becomes   saturated   with   chlorine,   no 
more  of  it  dissolves,  and  the  volume  of  gas  collected  at 
the  positive  pole  is  equal  to  that  of  the  hydrogen  at  the 
negative. 

103.  Occurrence  and  preparation  of  hydrochloric  acid. — 
Hydrochloric  acid  is  contained  in  the  gases  which  issue  from 
volcanoes.     Considerable  quantities  of  it  exist  dissolved  in 
the  waters  of  several  South  American  rivers,  whose  sources 
are  in  volcanic  districts  of  the  Andes.     Industrially  it  is 
always  obtained  by  the  action  of  sulphuric  acid,  H2S04,  on 
salt.     The  operation  is  carried  out  in  iron  pans  and  the 


CHLORINE 


81 


reaction  does  not  become  complete  until  the  mixture  is 
heated  nearly  to  redness : 

SNaCl  +  H2S04  :=  Na2S04  +  2HC1. 

104.  Preparation  of  hydrochloric  acid  gas  in  the  labora- 
tory.— For  laboratory  purposes  it  can  be  obtained  from  the 
same  materials.  But  it  is  more  convenient  to  start  with  the 
liquid  hydrochloric  acid  of  commerce.  This  is  a  solution 
of  the  gas,  HC1,  in  water,  made  by  passing  it  into  cold  water 
until  the  water  will  dissolve  no 
more.  Some  of  this  liquid  is 
placed  in  a  flask  with  three  necks. 
Through  one  passes  the  delivery 
tube,  which  conveys  the  gas  to  the 
jar  in  which  it  is  to  be  collected. 
Another  carries  a  dropping  funnel, 
from  which  strong  sulphuric  acid 
is  allowed  to  fall  drop  by  drop 
(Fig.  26).  We  shall  see  later  that 
sulphuric  acid  combines  energet- 
ically with  water.  Each  drop  com- 
bines with  a  certain  quantity  of  the 
water  in  the  hydrochloric-acid 
solution,  and  the  gas  which  was 
dissolved  in  this  escapes.  In  this 
way  an  abundant  supply  of  the  gas 
can  be  obtained.  When  the  stop- 
cock of  the  funnel  is  turned  so  that 
sulphuric  acid  no  longer  enters  the 
flask,  the  current  of  gas  ceases.  The  third  neck  carries  a 
safety-tube,  the  bend  of  which  contains  a  little  mercury. 
This  serves  to  prevent  the  pressure  from  rising  too  high  in 
the  bottle,  for  if  so  much  gas  is  liberated  that  the  delivery 
tube  is  unable  to  carry  off  all  of  it,  the  excess  will  escape 
through  the  safety-tube.  The  gas  can  not  be  collected  over 


PIG.  26.— Preparation  of  hydro- 
chloric-acid gas. 


82  ELEMENTARY  CHEMISTRY 

water  on  account  of  its  great  solubility.  Small  quantities 
can  be  collected  over  mercury,  on  which  it  has  no  action,  and 
larger  quantities  by  downward  displacement  in  dry  flasks. 

105.  Properties. — Hydrochloric  acid  is  a  colorless  gas, 
with  a  pungent,  irritating  odor.  By  cold  and  pressure  it 
has  been  converted  into  a  colorless  liquid,  which  has  been 
frozen  to  a  white  crystalline  solid.  Water  at  0°  dissolves 
500  times  its  bulk,  and  at  ordinary  temperatures  about  400 
times,  and  this  solution  is  the  hydrochloric,  or  "  muriatic," 
acid  of  commerce.  It  is  interesting  to  note,  that  while  the 
quantity  of  the  gas  dissolved  is  increased  by  increasing  the 
pressure,  yet  it  is  not  proportional  to  the  pressure,  as  we 
have  seen  to  be  the  case  with  slightly  soluble  gases  like 
oxygen.  More  gas  will  dissolve  under  two  atmospheres  than 
under  one,  but  nothing  like  twice  as  much.  All  very  soluble 
gases  behave  in  this  way. 

The  chemical  behavior  of  hydrochloric  acid  can  be 
summed  up  in  the  remark  that  it  is  a  strong  acid.  The 
term  "  strong  acid "  has  a  very  precise  meaning  in  mod- 
ern science,  and  we  shall  shortly  inquire  just  what  that 
meaning  is.  At  present  we  can  only  remark  that  hydro- 
chloric acid,  like  other  strong  acids,  is  energetic  chemically, 
readily  taking  part  in  many  chemical  changes ;  but  this  en- 
ergetic character  is  only  shown  when  it  is  dissolved — prac- 
tically we  may  say,  when  dissolved  in  water,  for  hydrochloric 
acid  free  from  water  is,  whether  gaseous  or  liquid,  an  inert 
substance. 

106.  Action  upon  the  metals. — Into  a  beaker  containing 
a  little  strong  aqueous  hydrochloric  acid  we  throw  a  small 
fragment  of  sodium.  The  metal  melts  and  runs  about  on 
the  surface  of  the  acid  with  a  hissing  noise.  Hydroger 
escapes  and  a  mass  of  fine  white  crystals  of  salt  falls  through 
the  liquid,  thus : 

HC1  +  Na  =  NaCl  +  H. 


CHLORINE  83 

The  aqueous  acid  acts  similarly  with  many  other  metals. 
When  we  studied  hydrogen  we  examined  its  behavior  with 
zinc.  We  can  now  write  the  equation, 

Zn  +  2HC1  =  ZnCl2  +  H2. 

So  with  magnesium,  there  is  energetic  evolution  of  hy- 
drogen— 

Mg  +  2HC1  =  MgCl2  +  H2. 

The  chloride  of  the  metal  used  is  formed  and  hydrogen 
escapes.  On  the  other  hand,  the  water-free  substance,  liquid 
or  gas,  has  no  action  upon  zinc,  magnesium,  or  sodium. 
Evidently  the  chlorides  can  be  looked  upon  as  hydrochloric 
acid  in  which  the  hydrogen  has  been  removed  and  a  metal 
inserted  in  its  place.  Thus,  23  parts  of  sodium  replace  1 
part  of  hydrogen,  producing  salt : 

HC1,  NaCl. 

But  65  parts  of  zinc  replace  2  parts  of  hydrogen,  producing 
zinc  chloride,  in  which  the  quantity  Zn  is  combined  with 
twice  35.5  parts  of  chlorine : 

or  ZnCl2. 

This  relation  between  hydrochloric  acid  and  the  chlo- 
rides of  the  metals  we  express  by  calling  the  metallic  chlo- 
rides the  salts  of  hydrochloric  acid.  And,  in  general,  a  salt 
of  an  acid  has  the  same  composition  as  the  acid  itself,  except 
that  the  hydrogen  has  been  replaced  by  a  metal. 

There  are  some  metals  on  which  the  action  of  hydro- 
chloric acid  is  slight,  silver,  for  instance;  and  others,  like 
gold  and  platinum,  on  which  it  does  not  act  at  all. 

107.  Decomposition  by  heat. — Hydrochloric  acid  is   a 
stable  compound  and  is  not  at  all  separated  into  hydrogen 
and  chlorine  by  a  temperature  of  1500°,  which  is  far  beyond 
a  white  heat.    At  1800°  it  is  partially  decomposed. 
7 


CHAPTER   XI 

THE  CHLORIDES—  COMPOUNDS   OF  CHLORINE   CONTAINING 
OXYGEN 

108.  Preparation  of  the  chlorides  of  the  metals.  —  Some 
metallic  chlorides  can  be  made  in  the  same  way  as  those  of 
zinc,  sodium,  and  magnesium  by  the  action  of  aqueous  hy- 
drochloric acid  upon  the  metal,  hydrogen  escaping.  There 
are  other  methods  of  obtaining  them.  To  a  strong  solution 
of  sodium  hydroxide  in  a  beaker  strong  aqueous  hydrochloric 
acid  is  added,  drop  by  drop,  with  constant  stirring.  There 
is  a  violent  reaction,  much  heat  is  evolved,  and  a  mass  of 
salt  crystals  separates: 

NaOH  +  HC1  =  STaCl  +  H20. 

Some  zinc  oxide  is  covered  with  water  and  hydrochloric 
acid  added,  with  stirring.  The  white  powder  disappears 
and  a  solution  of  zinc  chloride,  ZnCl2,  is  produced  : 

ZnO  +  2HC1  =  ZnCl2  +  H20. 

It  is  clear  from  the  equations,  that  if  the  materials  are 
pure  no  gas  will  be  given  off  in  either  experiment,  for  the 
hydrogen  does  not  escape.  It  forms  water.  The  efferves- 
cence which  usually  occurs  in  the  first  experiment  arises 
from  sodium  carbonate,  Na2C03,  which  is  nearly  always 
present  as  an  impurity  in  the  sodium  hydroxide.  We  have 
remarked  that  this  liberates  carbon  dioxide  with  hydro- 
chloric acid: 


Na2C03  +  2HC1  =  SNaCl  +  H20  +  C02. 
84 


THE  CHLORIDES  85 

This  leads  us  to  another  method  of  making  metallic 
chlorides.  The  carbonate  of  a  metal  is  treated  with  hydro- 
chloric acid.  Carbon  dioxide  escapes,  water  is  formed,  and 
the  chloride  in  question  dissolves  and  can  be  obtained  by 
evaporating  the  liquid.  This  is  the  method  usually  em- 
ployed in  making  calcium  chloride,  CaCl2.  Calcium  car- 
bonate, CaC03,  which  is  called  limestone,  marble,  or  chalk, 
according  to  its  condition,  is  dissolved  in  hydrochloric  acid: 

CaC03  +  2HC1  =  CaCl2  +  H20  +  CO,.1 

109.  Insoluble  chlorides. — The  chlorides  of  the  metals 
are  mostly  soluble  in  water.    Silver  chloride  (AgCl),  mer- 
curous  chloride  (Hg2Cl2),  and  cuprous  chloride   (Cu2Cl2) 
are  insoluble  in  it.2    Lead  chloride  (PbCl2)  is  slightly  solu- 
ble in  cold  water,,  much  more  so  when  the  water  is  hot. 
There  are  a  few  other  insoluble  and  slightly  soluble  chlo- 
rides of  less  importance. 

110.  Chlorides  of  non-metallic  elements. — The  chlorides 
of  non-metals  like  phosphorus  and  sulphur  can  not  be  ob- 
tained by  the  action  of  hydrochloric  acid  upon  the  element, 
for  the  acid  does  not  act  upon  non-metallic  elements  in  this 
way.    Nor  does  it  produce  chlorides  with  the  oxides  or  hy- 
droxides of  the  non-metals.     Only  the  metals  form  carbo- 
nates, so  that  the  third  method  can  not  be  applied.     This 
leaves  us  only  the  direct  action  of  chlorine  upon  the  ele- 
ment, and  it  is  in  this  way  that  the  non-metallic  chlorides 

1  These  three  methods  of  making  salts,  by  the  action  of  the  acid 
upon  the  oxide,  the  hydroxide,  or  the  carbonate  of  the  metal,  apply 
not  only  to  hydrochloric,  but  to  most  other  acids  as  well. 

2  To  call  a  substance  insoluble  is  simply  a  short  method  of  saying 
that  its  solubility  is  very  small.     Of  course,  instead  of  talking  in  this 
loose  way,  it  would  be  better  to  measure  the  solubility  of  these  sub 
stances  and  state  the  results  in  figures.     A  good  beginning  has  been 
made  in  this  direction.     Silver  chloride,  for  instance,  is  a  typical 
"insoluble"  substance,  yet  recently  its  solubility  has  not  only  been 
detected,  but  measured. 


86  ELEMENTARY  CHEMISTRY 

are  usually  prepared.  We  have  seen  the  method  applied  in 
the  case  of  phosphorus  and  arsenic.  While  the  chlorides  of 
the  metals  are  usually  soluble  in  water  unchanged,  those  of 
the  non-metals  react  chemically  with  it,  producing  hydro- 
chloric acid  and  other  products  whose  nature  depends  upon 
the  chloride  employed. 

COMPOUNDS  OF  CHLORINE  CONTAINING  OXYGEN 

111.  Three  oxides  of  chlorine  have  been  obtained: 

C120,  chlorine  monoxide, 
C102,  chlorine  peroxide, 
C1207,  chlorine  heptoxide. 

Chlorine  monoxide,  C120,  is  obtained  by  leading  chlorine 
over  cold  mercuric  oxide.  The  gas  given  off  is  condensed 
in  a  freezing  mixture.  It  is  yellowish-brown,  with  a  peculiar 
odor,  different  from  that  of  chlorine.  By  cooling  it  is  con- 
verted into  a  dark-brown  liquid.  This  can  be  distilled  with- 
out decomposition  if  pure  and  if  carefully  heated,  but  the 
operation  is  not  free  from  danger.  Both  the  liquid  and  the 
gas  explode  violently  on  being  heated  quickly,  and  on  con- 
tact with  many  substances — as,  for  instance,  phosphorus 
and  sulphur.  This  explosion  is  simply  decomposition  into 
chlorine  and  oxygen.  Since  the  oxygen  of  chlorine  monoxide 
is  only  loosely  held,  it  tends  to  yield  oxygen  to  any  sub- 
stance which  can  combine  with  that  element.  It  violently 
oxidizes  finely  divided  metals  and  many  other  substances, 
and  this  behavior  is  summed  up  in  the  statement  that  it 
is  a  strong  oxidizing  agent. 

112.  Chlorine  peroxide,  C102. — A  little  strong  sulphuric 
acid  is  placed  in  a  small,  strong  glass  cylinder,  and  0.5  gram 
or  less  of  finely  powdered  potassium  chlorate  introduced  in 
small  portions,  the  cylinder  being  covered  with  a  card  after 
each  addition.     Chlorine  peroxide,  a  dark,  greenish-yellow 
gas,  rises  and  fills  the  cylinder.     When  it  is  very  much 


COMPOUNDS  OF  CHLORINE  CONTAINING  OXYGEN  87 

diluted  with  air,  the  odor  of  the  gas  is  not  unpleasant. 
The  action  of  heat  upon  it  is  well  shown  by  bringing  into 
the  upper  part  of  the  cylinder  a  glass  rod  which  has  been 
heated  in  the  burner  flame.  Explosion  instantly  takes  place, 
the  gas  being  converted  into  chlorine  and  oxygen.  Its  be- 
havior on  contact  with  combustible  substances  can  be  illus- 
trated by  allowing  a  drop  of  ether  to  fall  into  another  jar 
of  the  gas.  There  is  an  explosion.  The  ether  is  burned 
to  carbon  dioxide  and  water  at  the  expense  of  the  oxygen  of 
the  chlorine  peroxide.  Larger  quantities  of  chlorine  perox- 
ide can  be  made  by  gently  heating  potassium  chlorate  with 
strong  sulphuric  acid  in  a  flask,  but  the  operation  is  dan- 
gerous. When  led  into  a  tube  surrounded  by  a  freezing 
mixture  of  ice  and  salt,  the  gas  condenses  to  a  red  liquid, 
which,  on  more  intense  cooling,  solidifies  to  a  mass  of  yel- 
low crystals. 

113.  Chloric  acid,  HC103,  has  never  been  obtained  free 
from  water.    The  strong  solution  is  thick  and  colorless,  and 
tends   energetically  to   impart   oxygen  to   oxidizable   sub- 
stances; paper  wet  with  it  catches  fire  after  a  time.    Corre- 
sponding to  chloric  acid  is  a  series  of  compounds  in  which 
its  hydrogen  is  replaced  by  metals,  e.  g. : 

KC103,  potassium  chlorate, 
NaC103,  sodium  chlorate. 

These  compounds  are  called  the  chlorates. 

114.  Potassium  chlorate,  KC103,  was  the  first  to  be  dis- 
covered (1786),  and  is  still  the  most  important.    It  can  be 
obtained  by  the  action  of  chlorine  on  hot  strong  solution  of 
potassium  hydroxide: 

6KOH  +  3C12  =  5KC1  +  KC103  +  3H20. 

Since  f  of  the  potassium  present  is  wasted  as  potassium 
chloride  in  this  reaction,  the  process  is  no  longer  used.  Po- 
tassium chlorate  is  now  made  almost  entirely  by  the  action 


88  ELEMENTARY   CHEMISTRY 

of  the  electric  current  upon  a  solution  containing  potassium 
chloride,  KC1,  and  potassium  hydroxide,  KOH.  It  would 
lead  us  too  far  to  discuss  the  details  of  this  interesting 
process  here.  The  potassium  chlorate  separates  out  in  crys- 
tals, which  can  be  removed  with  a  perforated  ladle  while 
fresh  potassium  chloride  is  added,  so  that  the  process  is  con- 
tinuous. Potassium  chlorate  forms  white  anhydrous  crys- 
tals, which  are  only  slightly  soluble  in  cold  water,  much 
more  so  in  hot.  Its  solution  is  used  in  medicine  as  a  gargle 
for  inflamed  mucous  membranes  of  the  throat.  It  is  poison- 
ous, and  must  not  be  swallowed. 

The  torpedoes  which  are  placed  on  railway  tracks  as 
signals  consist  of  small  tin  boxes,  like  blacking  boxes,  filled 
with  a  mixture  of  powdered  sulphur  and  potassium  chlorate ; 
and  we  can  understand  their  behavior  by  grinding  a  small 
crystal  of  potassium  chlorate  with  a  fragment  of  sulphur  in 
a  mortar.  The  sulphur  is  oxidized  with  explosion  by  the 
oxygen  of  the  potassium  chlorate.  This  illustrates  the  fact 
that  mixtures  of  chlorates  with  oxidizable  materials  are 
explosive,  and  such  mixtures  have  been  used  practically  for 
blasting,  but  they  are  now  abandoned,  because  they  are  sen- 
sitive to  shock,  and  it  is  impossible  to  transport  them  safely. 

Chlorine  heptoxide,  C1207,  is  a  colorless  oil  which  is 
much  more  stable  than  the  other  oxides  of  chlorine.  It  ex- 
plodes on  being  brought  into  a  flame  or  when  sharply  struck, 
but  it  has  no  action  on  wood,  paper,  phosphorus,  or  other 
oxidizable  materials. 

115.  Bleaching  powder,  CaOCl2. — Some  slaked  lime, 
which  is  calcium  hydroxide,  Ca(OH)2,  is  mixed  to  a  paste 
with  water,  and  poured  into  a  flask  filled  with  chlorine. 
Then  the  flask  is  closed  with  the  hand  and  shaken  vigorously. 
The  color  of  the  chlorine  disappears,  and  the  suction  upon 
the  hand — the  flask  will  hang  suspended  from  it — shows 
that  the  gas  is  being  absorbed.  When  the  absorption  is  com- 
plete we  add  a  little  nitric  acid  to  the  white  substance  in 


COMPOUNDS  OF  CHLORINE  CONTAINING  OXYGEN  89 

the  flask.  The  chlorine  is  again  liberated,  and  soon  fills  the 
vessel. 

Bleaching  powder,  or  "  chloride  of  lime,"  as  it  is  called, 
is  made  in  large  stone  chambers.  On  the  floor  a  thin  layer 
of  slaked  lime  is  placed  and  chlorine  is  passed  in  until  the 
chamber  is  full  of  the  gas.  Then  the  chamber  is  closed  and 
allowed  to  stand  two  or  three  days,  after  which  the  lime  is 
raked  up,  exposing  fresh  surfaces,  and  the  chlorine  treat- 
ment repeated. 

With  respect  to  the  chemical  make-up  of  bleaching  pow- 
der there  has  been  much  discussion.  Most  of  the  facts  are 
fairly  well  summed  up  in  the  statement  that  it  consists 
essentially  of  a  compound  CaOCl2,  which  liberates  its  chlo- 
rine when  treated  with  an  acid — e.  g. : 

CaOCl2  +  H2S04  =  CaS04  +  H20  +  C12. 

When  well  made  it  liberates  40  per  cent  or  upward  of  chlo- 
rine in  this  way.  It  is  simply  a  convenient  way  of  preserv- 
ing chlorine  and  shipping  it,  and  all  its  uses  depend  upon  the 
ease  with  which  chlorine  can  be  obtained  from  it.  It  is  used 
as  a  disinfectant  and,  in  enormous  quantities,  for  bleaching. 


CHAPTEB   XII 

THE  ATOMIC  THEORY-THE  LA  W  OF  MULTIPLE  PROPORTIONS 

116.  The  subdivision  of  chalk. — Ordinary  chalk  is  a 
white,  odorless,  tasteless,  brittle  solid.  If  a  piece  of  it  is 
broken  in  two,  each  portion  remains  chalk  as  before — there 
is  no  alteration  in  properties  as  a  result  of  the  division. 
Instead  of  breaking  the  chalk  with  the  fingers,  it  will  be 
easier  to  put  it  in  a  mortar  and  grind  it.  In  this  way  we 
can  pulverize  chalk  very  finely,  but  we  soon  reach  the  limit 
of  the  process — the  powder  packs  together  in  the  bottom  of 
the  mortar,  and  further  grinding  is  useless.  If,  now,  we 
examine  some  of  this  powder  under  a  microscope,  it  be- 
comes clear  that  every  grain  of  it  is  simply  a  little  mass 
of  chalk,  which  possesses  a  definite  shape  and  has,  in  gen- 
eral, all  the  properties  which  we  commonly  associate  with 
that  substance. 

Here  we  are  at  the  end  of  our  resources.  So  far  as 
we  know  practically,  chalk  remains  chalk  no  matter  how 
fine  the  grains  of  the  powder  become. 

Now  let  us  take  another,  a  very  important,  step.  Let  us 
carry  forward  the  division  of  the  chalk  in  thought.  We  can 
pick  out  a  grain  of  the  powder  and  imagine  a  plane  passed 
through  it  so  as  to  divide  it  into  two  equal  parts.  Then 
repeat  this  operation  with  each  half,  and  so  on.  Clearly 
this  is  purely  a  mental  process,  and  the  results  we  get  from 
it  will  be  ideas,  not  facts. 

At  once  we  are  face  to  face  with  two  possibilities,  both 
of  which  are  as  old  as  human  knowledge  itself : 
90 


THE  ATOMIC  THEORY  91 

First.  The  subdivision  of  the  chalk  can  be  carried  out 
without  limit.  No  matter  how  small  a  mass  of  chalk  we 
think  of,  it  can  still  be  subdivided  into  two  smaller  masses. 
This  view,  when  applied  not  only  to  chalk  but  to  all  matter, 
is  called  the  doctrine  of  the  infinite  divisibility  of  matter. 

Second.  Ultimately  we  should  reach  a  particle  of  chalk 
so  small  that  if  a  plane  were  passed  through  it,  different 
kinds  of  matter  would  lie  on  the  two  sides  of  the  plane. 
The  chalk  would  cease  to  be  chalk  when  this  subdivision 
was  made,  and  would  yield  two  new  substances.  This  view 
is  the  opposite  of  the  preceding  one.  It  may  be  called  the 
doctrine  that  matter  is  not  infinitely  divisible,  or,  more 
shortly,  the  atomic  theory. 

So  far  as  the  subdivision  of  chalk  is  concerned,  we  are 
not  forced  to  choose  between  these  alternatives.  We  might 
answer  that,  not  being  able  to  carry  out  the  process,  we  do 
not  know  whether  chalk  is  infinitely  divisible  or  not.  Never- 
theless, for  practical  reasons,  it  is  necessary  to  make  a 
choice.  It  has  been  found  that  the  second  mew,  the  atomic 
theory,  aids  us  immensely  in  remembering  and  classifying 
physical  and  chemical  facts,  while  the  first  is  nearly  worth- 
less in  these  respects. 

117.  Molecules, — The  smallest  mass  of  chalk  which  we 
can  think  of — that  mass  which  ceases  to  be  chalk  and  be- 
comes something  else  when  we  divide  it  further — we  shall 
call  a  molecule  of  chalk.     Any  mass  of  chalk  consists  of 
molecules,  all  of  which  are  alike.     So  does  any  other  pure 
substance,  but  the  molecules  of  two  different  substances  are 
different.     These  molecules  are  not  to  be  thought  of  as 
being  in  contact ;  there  are  spaces  between  them,  and,  if  we 
like,  we  may  think  of  these  spaces  as  filled  with  the  ether 
of  physics,  though  this  assumption  does  not  help  us  much 
in  dealing  with  chemical  facts. 

118.  Movement  of  the  molecules  in  solids,  liquids,  and 
gases. — In  a  solid,  a  molecule  moves  constantly,  but,  on  the 


92  ELEMENTARY  CHEMISTRY 

whole,  remains  in  about  the  same  locality,  the  motion  being 
like  that  of  the  earth  around  the  sun  or  of  a  vibrating  tun- 
ing-fork. In  a  liquid  or  a  gas,  a  molecule  moves  in  a 
straight  line  until  it  strikes  another  molecule  or  the  wall 
of  the  containing  vessel,  the  motion  being  like  that  of  a 
swarm  of  bees  enclosed  in  a  box.  When  one  of  these  en- 
counters occurs  the  molecule  does  not  ordinarily  return  to 
the  same  place,  but  rebounds  in  some  new  direction.  The 
great  difference  between  liquids  and  gases  is,  that  in  a  liquid 
the  molecules  are  crowded  together,  and  each  one  only  trav- 
els a  very  short  distance  before  it  strikes  another ;  in  a  gas, 
the  crowding  is  not  so  great,  and  the  free  path  of  a  mole- 
cule— the  distance  it  travels  before  an  encounter  takes  place 
— is  longer. 

These  distinctions  are  only  rough.  Many  facts — for  in- 
stance, the  traveling  of  gold  into  lead  (p.  19) — show 
that,  under  some  circumstances,  molecules  of  solids  must 
be  assumed  to  desert  their  original  position  and  move 
forward  to  a  new  one  some  distance  away,  and  mole- 
cules continually  leave  the  surface  of  all  liquids  and  many 
solids  and  begin  to  play  the  role  of  gaseous  molecules  in 
the  surrounding  space.  This  occurrence  is  evaporation, 
and  the  molecular  motion  of  solids,  liquids,  and  gases  is 
heat. 

119.  Size  of  the  molecules. — Some  physical  facts  are  best 
dealt  with  by  giving  a  definite  size  to  the  molecule.  Ac- 
cording to  these  facts,  about  5,000,000  molecules  laid  side 
by  side  in  contact  would  make  a  line  one  centimeter  long. 
This  corresponds  to  about  12,500,000  of  them  in  an  inch. 
In  this  calculation  it  is  assumed  that  the  molecules  are 
spheres.  This  shape  the  chemist  can  not  give  to  the  mole- 
cules by  means  of  which  he  describes  chemical  phenomena, 
because  his  facts  will  not  allow  him  to  do  so. 

At  present,  chemical  facts  do  not  force  us  to  assign  any 
size  to  the  molecules.  It  is  sufficient  to  think  of  them  as 


THE  ATOMIC  THEORY  93 

very  minute — so  small  that  the  smallest  visible  fragment  of 
any  substance  must  be  thought  of  as  containing  many  mil- 
lions. 

120.  Atoms. — It  has  been  said  that,  from  the  standpoint 
of  the  atomic  idea,  any  mass  of  chalk  is  to  be  thought  of 
as  consisting  of  molecules,  all  of  which  are  exactly  alike. 
Let  us  now  place  some  chalk  in  a  vessel  which  will  stand  a 
very  high  temperature — a  platinum  vessel  is  best — and  heat 
it  to  whiteness,  arranging  the  apparatus  in  such  a  way  that 
any  gas  given  off  can  be  collected  over  mercury.  A  color- 
less gas  collects  over  the  mercury,  which  can  be  shown  to 
be  carbon  dioxide,  C02.  When  the  operation  is  over,  we 
open  the  platinum  vessel  and  find  in  it  common  lime,  which 
is  calcium  oxide,  CaO. 

Now,  since  every  molecule  of  chalk  is  like  every  other, 
each  one  must  have  separated  into  at  least  one  molecule  of 
lime  and  one  of  carbon  dioxide.  By  methods  which  we  need 
not  discuss,  lime  can  be  proved  to  contain  calcium  and  oxy- 
gen and  carbon  dioxide  to  contain  carbon  and  oxygen,  and  it 
is  clear  that  each  molecule  of  chalk  must  be  thought  of  as 
containing  these  three  constituents. 

Hence  the  necessity  for  thinking  of  the  molecule,  not  as 
a  simple  mass,  but  as  a  structure  built  up  of  smaller  masses, 
the  ATOMS.  We  must  think  of  each  molecule  of  chalk  as 
containing  at  least  four  atoms,  one  of  calcium,  one  of  car- 
bon, and  two  of  oxygen,  one  of  which  goes  with  the 
calcium  atom  and  the  other  with  the  carbon  atom  when 
the  molecule  is  broken  up  by  heat.  By  further  reasoning 
it  can  be  shown  to  contain  five  atoms,  for  the  molecule  of 
carbon  dioxide  contains  two  atoms  of  oxygen,  both  of  which 
must  have  come  from  the  chalk,  and  therefore  there  must 
be  three  oxygen  atoms  present,  not  two. 

Atoms,  then,  are  the  smaller  particles  of  which  mole- 
cules consist.  The  atoms  of  the  same  element  are  all 
exactly  alike;  every  atom  of  hydrogen  is  like  every  other 


94  ELEMENTARY  CHEMISTRY 

atom  of  hydrogen;  every  oxygen  atom  is  like  every  other. 
When  we  say  that  the  atomic  weight  of  oxygen  is  16,  then, 
if  we  happen  to  be  thinking  of  molecules  and  atoms  and  not 
of  matter  in  visible  masses,  what  we  mean  is  this :  that  the 
oxygen  atom  weighs  16  times  as  much  as  the  hydrogen  atom. 
It  makes  no  difference  what  the  weight  of  the  hydrogen 
atom  is.  That  does  not  interest  us  in  the  least.  We  are 
concerned  only  with  the  statement  that,  whatever  the  weight 
of  the  hydrogen  atom  may  be,  that  of  the  oxygen  atom  is 
16  times  as  great. 

Since  all  hydrogen  atoms  are  alike  and  all  oxygen  atoms 
are  alike,  this  is  a  perfectly  definite  statement,  and  is  true 
of  all  pure  hydrogen  and  oxygen  without  regard  to  the 
source  or  the  method  of  preparation. 

121.  The  chemical  laws  from  the  standpoint  of  atoms 
and  molecules. — It  will  now  interest  us  to  go  over  the  chem- 
ical laws,  which  we  have  previously  studied  simply  as  facts, 
and  to  see  what  appearance  they  present  when  we  make 
use  of  the  ideas  of  molecule  and  atom  in  stating  them.    In 
doing  this  we  are  not  bringing  forward  anything  new.    We 
are  simply  stating  the  same  old  facts  in  a  new  language — 
the  language  of  the  atomic  theory. 

122.  The  law  "of  the  indestructibility  of  matter. — In  the 
first  place,  the  atoms  are  not  destroyed  and  are  not  created 
in  chemical  operations,  they  simply  continue  to  exist.     As 
John  Dalton,  the  founder  of  the  atomic  theory  in  its  pres- 
ent form,  remarked  in  1808 :  "  We  might  as  well  attempt 
to  introduce  a  new  planet  into  the  solar  system,  or  to  anni- 
hilate one  already  in  existence,  as  to  create  or  destroy  a 
particle  of  hydrogen."    Hence  the  law  of  the  indestructibil- 
ity of  matter.     Each  atom  is  conceived  as  indestructible, 
and,  since  all  matter  consists  of  atoms,  no  destruction  is 
possible.    After  any  chemical  change  the  number  of  atoms 
remains  the  same  as  before  it. 

On  the  other  hand,  the  total  number  of  molecules  can 


THE  ATOMIC  THEORY  95 

be  very  greatly  altered  by  chemical  changes.  We  have  just 
discussed  the  decomposition  of  chalk  by  heat.  The  equation 
for  this  decomposition  is — 

CaC03  =  CaO  +  C02, 

and  if  we  consider  that  each  symbol  means  an  atom  of  the 
corresponding  element  and  each  formula  a  molecule  of  the 
compound,  we  perceive  that  every  molecule  of  chalk  yields 
two  molecules,  one  of  lime  and  one  of  carbon  dioxide,  so  that 
the  total  number  of  molecules  present  is  doubled.  But  the 
number  of  atoms  remains  five  as  before. 

123.  The  law  of  definite  proportions. — Similarly  with 
the  law  of  definite  proportions.     The  formula  of  hydro- 
chloric acid  is  HC1,  and  we  have  thought  of  this  as  meaning 
one  atomic  weight  of  hydrogen,  weighing  1  (1  gram,  for  in- 
stance), and  one  atomic  weight  of  chlorine,  weighing  35.5 
(35.5  grams).     But  we  are  at  liberty  to  think  of  any  unit 
of  weight  we  please.     Let  us  make  our  unit  of  weight  the 
hydrogen  atom.    Then  the  formula  HC1  means  an  atom  of 
hydrogen,  weighing  1,  in  union  with  an  atom  of  chlorine, 
weighing  35.5  times  as  much.    These  together  form  a  mole- 
cule of  .hydrochloric  acid,  which  must  weigh  36.5  times  as 
much  as  the  hydrogen  atom. 

Both  these  atoms  possess  an  invariable  weight,  and 
therefore  the  composition  of  a  molecule  of  hydrochloric  acid 
must  be  always  the  same.  It  must  contain  ^  of  its 
weight  of  hydrogen  and  |^|  of  chlorine.  But  any  mass  of 
hydrochloric  acid  consists  simply  of  a  great  number  of  such 
molecules,  and  its  composition  must  be  the  same  as  that  of 
a  single  molecule.  Accordingly,  the  composition  of  hydro- 
chloric acid  must  be  constant.  The  same  reasoning  applies 
to  all  compounds. 

124.  The  law  that  it  is  impossible  to  transform  one  ele- 
ment into  another. — Since  an  atom   simply  continues  to 
exist  and  can  not  be  converted  into  any  other  atom,  the 


96  ELEMENTARY   CHEMISTRY 

atomic  theory  gives  us  a  clear  picture  of  the  fact  that  one 
element  is  never  transformed  into  another. 

125.  The  atomic  weights. — We  have  seen  that  there  is 
for  each  element  a  natural  quantity  by  weight  in  which  it 
enters  into  its  compounds.    According  to  the  atomic  theory, 
these  numbers  are  simply  the  relative  weights  of  the  atoms, 
that  of  the  hydrogen  atom  being  1.    The  molecule  of  hydro- 
chloric acid  contains  two  atoms,  one  of  hydrogen  and  one 
of  chlorine,  and  the  chlorine  atom  weighs  35.5  times  as  much 
as  the  hydrogen  atom.     Every  compound  of  chlorine  con- 
sists of  molecules  which  contain  chlorine  atoms  in  union 
with  those  of  other  elements.     Thus,  every  compound  of 
chlorine  must  contain  35.5  parts  of  chlorine  by  weight  or — 
if  it  contains  more  than  one  atom  of  chlorine — some  mul- 
tiple of  that  quantity. 

126.  The  law  of  multiple  proportions. — The  two  com- 
pounds of  hydrogen  and  oxygen,  water  (H20),  and  hydro- 
gen peroxide  (H202),  have  been  described.     Starting  with 
water,  let  us  admit  that  there  is  a  second  compound  of  the 
same  two  elements,  richer  in  oxygen  but  containing  the 
same  two  atoms  of  hydrogen  in  the  molecule.     Then  it  is 
clear  that  this  second  compound  must  contain  at  least  two 
atoms  of  oxygen,  and  have  the  formula  H202,  for  we  can  not 
divide  the  oxygen  atom,  and  any  compound  between  the  two 
is  impossible.     Still  increasing  the  oxygen,  the  next  possi- 
bility would  be  H203  and  the  next  H204,  and  there  are  in- 
dications that  these  compounds  exist,  though  they  have  not 
yet  been  obtained. 

Statements  precisely  like  this  can  be  made  for  every 
case  where  the  same  two  elements  form  more  than  one 
compound.  In  every  such  case,  if  we  suppose  the  quantity 
of  one  element  to  remain  the  same,  the  quantity  of  the 
other  element  will  increase  by  a  leap,  so  that  its  quantity 
in  the  second  compound  is  double  or  triple  or  some  small 
multiple  of  that  in  the  first.  Intermediate  compounds  do 


THE  LAW  OP  MULTIPLE  PROPORTIONS  97 

not  exist.  This  fact  is  commonly  called  the  law  of  multiple 
proportions.  As  we  have  just  seen,  our  atomic  theory  gives 
us  a  simple  account  of  it.  For,  if  we  suppose  the  quantity 
of  the  second  element  to  increase  at  all,  it  must  increase 
by  at  least  one  atom  in  the  molecule.  And  since  the  atoms 
are  not  divided,  no  compound  between  can  exist. 


ciwjuwtjrt  43  \Af4(A&L  uwdju  wi)d.  AS! 
(^/  /  if 

/TT^~t^~T*  £ju^  $ihA) . 


CHAPTER   XIII 

THE  ATMOSPHERE— NITROGEN 

127.  Historical. — Aristotle,  whose  views  were  accepted 
without  question  through  the  middle  ages,  considered  the 
air  to  be  an  element.  During  the  seventeenth  century  Boyle 
opposed  this  idea,  and  announced  that  it  was  probably  a 
complicated  mixture.  This  opinion  rapidly  gained  ground 
during  the  hundred  years  that  followed,  but  it  was  not  sup- 
ported by  convincing  experiments  until  near  the  close  of 
the  eighteenth  century  (1772).  We  already  know  that  when 
a  candle  is  burned,  the  products  are  the  gas  carbon  dioxide 
and  the  vapor  of  water.  If  the  candle  is  placed  in  a  closed 
vessel,  it  will  go  out  when  the  oxygen  is  all,  or  nearly  all, 
converted  into  these  two  substances.  Then,  when  the  vessel 
cools,  the  water-vapor  will  condense  on  its  walls  and  the 
gas  in  it  is  a  mixture  of  carbon  dioxide  with  air  from  which 
the  oxygen  has  been  removed,  for  the  other  constituents 
are  not  affected  by  the  combustion.  The  next  step  is  to  get 
rid  of  the  carbon  dioxide,  and  this  is  easily  accomplished 
by  introducing  a  little  lime,  which  greedily  absorbs  it. 
Daniel  Rutherford,  an  English  physician,  carried  out  this 
experiment  in  1772,  and  it  is  clear  that  it  amounts  to  simply 
subtracting  the  oxygen  from  the  air.  He  found  the  residual 
gas  to  be  colorless  and  odorless.  It  extinguished  burning 
substances  and  he  named  it  mephitic  air,  because  it  instant- 
ly suffocated  a  mouse  placed  in  it.  By  allowing  a  mouse  to 
suffocate  in  a  sealed  vessel  full  of  air,  and  then  absorbing 
the  carbon  dioxide  which  its  lungs  had  produced  by  lime,  he 
98 


DANIEL  EUTHEEFOED 
B.  Edinburgh,  1749.     D.  1819. 


THE  ATMOSPHERE— NITROGEN  99 

obtained  the  same  gas.  At  the  time,  and  for  more  than  a 
century  afterward,  it  was  considered  an  element  and  named 
nitrogen.  Within  the  last  few  years  it  has  been  found  that, 
while  the  gas  obtained  by  removing  the  oxygen  from  air  does 
consist  very  largely  of  nitrogen,  it  also  contains  small  quan- 
tities of  argon  and  other  elements  recently  discovered. 

128.  Composition  of  the  atmosphere. — The  air  consists 
mainly  of  the  two  gases,  nitrogen  and  oxygen,  and  if  every- 
thing else  were  removed  from  it  it  would  contain  by  vol- 
ume— 

Nitrogen,  79  per  cent ; 
Oxygen,  21  per  cent; 
by  weight — 

Nitrogen,  77  per  cent; 
Oxygen,  23  per  cent. 

129.  The  air  is  a  mixture,  not  a  compound. — These  gases 
are  not  chemically  united,  they  are  simply  mixed.    Perhaps 
the  strongest  proof  of  this  statement  is  the  fact  that  the 
percentage  of  oxygen  in  the  air,  .even  in  pure  open  air  of 
the  country  or  seashore,  varies.    Sometimes  it  falls  as  low 
as  20.8  per  cent,  while  in  the  air  of  mines  and  crowded 
rooms  it  may  drop  to  20.2  per  cent.    If  the  air  were  a  com- 
pound, the  proportion  of  oxygen  in  it  would  be  always  the 
same. 

130.  Analysis  of  air. — There  are  various  ways  of  ana- 
lyzing the  air.    One  is  to  pass  air — which  must  be  carefully 
purified  beforehand — through  a  weighed  tube  containing 
red-hot  copper.     This  absorbs  the  oxygen,  producing  black 
cupric  oxide  (CuO).     The  increase  in  weight  of  this  tube 
is  the  quantity  of  oxygen.     The  rest  of  the  air  (nitrogen 
mainly,  but  containing  also  argon  and  traces  of  other  ele- 
ments) passes  into  a  glass  globe  which  has  been  previously 
exhausted  at  an  air-pump  and  weighed.     The  increase  in 
weight  of  the  globe  gives  the  quantity  of  nitrogen  with 

8 


100  ELEMENTARY  CHEMISTRY 

argon  and  the  other  elements  which  are  present  only  in 
small  quantities. 

131.  Solubility  of  air  in  water. — Both  nitrogen  and  oxy- 
gen dissolve  in  water,  and  when  we  expel  air  from  water 
by  boiling  and  analyze  it,  we  find  that  it  contains  35  per  cent 
of   oxygen,   considerably   more   than   ordinary   air.      This 
difference  is  very  important  to  marine  life.     At  the  same 
time  it  furnishes  an  interesting  proof  that  the  air  is   a 
mixture,  for  if  it  were  a   compound  the  composition  of 
the  dissolved  air  would  be  the  same  as  that  of  the  origi- 
nal air. 

132.  Other  constituents  of  the  air. — The  nitrogen  ob- 
tained by  removing  oxygen  from  air  contains  a  little  more 
than  1  per  cent  by  volume  of  argon,  and  small  quantities 
of  four  other  elements — helium.,  krypton,  neon,  and  xenon — 
about  which  little  is  known  at  present,  but  which  appear 
to  resemble  argon  closely.    All  these  elements  are  colorless 
gases.    The  most  interesting  peculiarity  of  argon,  and  prob- 
ably of  the  others,  is  its  complete  inertness  chemically.1 
All  attempts  to  induce  it  to  take  part  in  chemical  changes 
have  failed,  and  it  appears  that  we  have  here  the  first  ex- 
ample of  an  element  altogether  destitute  of  chemical  activ- 
ity, a  new  form  of  matter. 

Water-vapor  is  always  present  in  the  air,  the  quantity 
being  very  different  at  different  times.  It  is  called  "  hu- 
midity" in  the  weather  reports,  and  much  of  it  makes  a 
moderately  warm  day  very  oppressive  because  it  interferes 
with  the  evaporation  from  the  body.  The  air  contains  about 

1  Helium  is  just  as  inert  as  argon ;  the  others  have  not  yet  been 
investigated  in  this  respect.  Helium  is  the  only  gas  which  has  thus 
far  resisted  all  attempts  to  liquefy  it.  It  has  been  strongly  compressed 
and  cooled  to  —260°  by  means  of  boiling  liquid  hydrogen  without 
liquefaction.  If  the  helium  is  then  allowed  to  expand  suddenly,  a 
further  fall  of  temperature  to  —204°  occurs,  and  the  gas  appears 
misty  from  incipient  condensation,  but  no  visible  liquid  is  produced. 


THE  ATMOSPHERE— NITROGEN 


101 


3  parts  in  10,000  by  volume  of  carbon  dioxide  (C02).  The 
quantity  varies  but  little;  it  rises  slightly  in  winter  in  the 
temperate  zone,  because  plant  life  is  mostly  at  a  standstill 
and  combustion  and  respiration  are  more  active.  On  this 
small  percentage  of  carbon  dioxide  the  existence  of  the  whole 
vegetable  world,  and  through  it  of  animals  also,  depends. 
Eecently  the  existence  of  a  little  hydrogen  (2  parts  in 
10,000)  has  been  announced.  There  are  present  also  traces 
of  ammonia  (NH3)  and  nitric  acid  (HN03),  which  are 
washed  down  into  the  soil  by  rains  and  furnish  an  impor- 
tant source  of  nitrogen  to  plants. 

NITROGEN,  N"  =  14 

133.  The  discovery  of  nitrogen  and  its  occurrence  in 
the  air  have  already  been  discussed.    It  is  not  an  important 
constituent  of  the  earth's  crust.   The 

chief  compounds  of  the  element 
which  occur  in  nature  are  potassium 
nitrate,  KN03,  commonly  called  salt- 
peter, and  sodium  nitrate,  NaN03, 
and  these  are  found  only  in  small 
quantities  except  in  some  special  lo- 
calities. 

134.  Preparation. — It  is  easy  to 
remove  the  oxygen  from  air,  and  the 
gas  which  remains — although,  as  we 
have    seen,   it    still   contains    argon 
and     several     other     elements — be- 
haves  like   pure   nitrogen   in   most 
respects.    We  can  pass  air  through 
a   tube   containing   red-hot    copper, 
which  will  retain  the  oxygen  in  the 
form  of  cupric  oxide.    The  nitrogen 

passes  on  unchanged,  and  can  be  collected  over  water.  Or 
a  graduated  tube  full  of  air  and  sealed  at  one  end  is  placed 


FIG.  27.— Analysis  of  air  by 
means  of  phosphorus. 


102  ELEMENTARY  CHEMISTRY 

with  the  open  end  in  water,  and  a  piece  of  phosphorus 
thrust  up  into  the  air  on  a  wire  (Fig.  27).  Phosphorus 
combines  slowly  with  oxygen  at  ordinary  temperatures,  and 
in  the  course  of  several  days  it  will  completely  remove  the 
oxygen  of  the  air  in  the  tube.  This  gradual  absorption  is 
shown  by  a  gradual  rise  of  the  water  into  the  tube.  About 
one-fifth  of  the  air  is  absorbed.  Then  the  level  of  the  water 
remains  constant  because  nitrogen  and  the  other  elements 
are  not  acted  upon  by  phosphorus. 

In  order  to  obtain  nitrogen  free  from  argon  we  must 
resort  to  chemical  methods  of  making  it.  One  way  is  to 
pass  chlorine  into  an  aqueous  solution  of  the  compound  of 
nitrogen  and  hydrogen — ammonia,  NH3.  The  chlorine  robs 
the  ammonia  of  its  hydrogen,  liberating  free  nitrogen: 

NH3  +  3C1  =  3HC1  +  N. 

The  hydrochloric  acid  does  not  remain  free.  It  unites  with 
more  ammonia,  producing  ammonium  chloride,  NH4C1: 

NH3  +  HC1  =  NH4C1, 

so  that  the  final  results  are  the  liberation  of  nitrogen  as 
gas  and  the  production  of  ammonium  chloride  in  the  liquid. 
This  experiment  must  be  carried  out  with  care,  for  if  we 
pass  in  too  much  chlorine  there  is  produced  nitrogen  chlo- 
ride, NC13,  a  dangerously  explosive  substance. 

A  better  method  is  to  heat  a  white  solid  called  ammoni- 
um nitrite,  which  has  the  composition  NH4N02.    On  gentle 
heating  it  separates  into  water  and  nitrogen: 
NH4N02  =  2H20  +  £T2. 

It  is  somewhat  difficult  to  obtain  ammonium  nitrite  pure, 
and  for  this  purpose  quite  unnecessary.  Practically,  we  use 
a  mixture  of  sodium  nitrite,  NaN02,  with  ammonium 
chloride,  NH4C1.  The  first  thing  that  happens  is  the  pro- 
duction of  ammonium  nitrite: 

NH4C1  =  NH4N08  +  NaCl. 


THE  ATMOSPHERE— NITROGEN  103 

This  then  decomposes : 

NH4N02  =  2H20  +  Na. 

The  mixture  is  placed,  with  enough  water  to  form  a  thin 
paste,  in  a  small  flask  provided  with  a  delivery  tube,  and 
carefully  heated.  The 
nitrogen  is  collected 
over  water  (Fig.  28). 

135.  Physical    prop- 
erties.— Nitrogen    is    a 
colorless,    odorless    gas. 
Its   solubility   in   water 
is  very  slight,  100  vol- 
umes   Of   the    latter    dlS-    FlG-  ^--Preparation  of  nitrogen  from  ammo- 

.    .  nium  nitrite. 

solving  less  than  1|  vol- 
umes of  the  gas  at  ordinary  temperatures.  From  this 
we  should  expect  it  to  be  difficult  to  liquefy,  and  this  is  the 
case;  yet  by  pressure  and  cold  together  it  has  been  con- 
verted into  a  colorless  liquid.  When  liquid  nitrogen  evapo- 
rates it  absorbs  heat  rapidly  from  surrounding  objects,  pro- 
ducing intense  cold ;  and  if  the  evaporation  is  made  more 
rapid  by  placing  the  vessel  under  an  air-pump,  the  tempera- 
ture falls  to  — 225°  and  the  nitrogen  freezes  to  a  snow-like 
mass.  This  was  for  a  long  time  the  greatest  cold  attain- 
able, but  recently  temperatures  more  than  30  degrees  lower 
have  been  obtained  by  the  evaporation  of  liquid  hydrogen. 

Since  the  air  is  nearly  four-fifths  nitrogen  by  volume, 
it  is  hardly  necessary  to  remark  that  the  gas  is  not  poison- 
ous. On  the  other  hand,  it  is  clear  that  an  animal  placed 
in  it  must  immediately  die  by  suffocation,  which  is  only 
another  name  for  lack  of  oxygen. 

136.  Chemical  properties. — Chemically,  nitrogen  is  in- 
ert at  ordinary  temperatures.    Burning  substances — a  can- 
dle, phosphorus,  and  charcoal,  for  instance — plunged  in  it 
are  extinguished  at  once.     At  higher  temperatures  it  is 


104  ELEMENTARY  CHEMISTRY 

more  energetic,  and  combines  directly  with  many  other  ele- 
ments— e.  g.,  boron,  lithium,  and  magnesium — producing 
with  them  compounds  called  nitrides;  for  example,  magne- 
sium nitride,  Mg3N2.  When  a  stream  of  electric  sparks  is 
passed  through  a  mixture  of  nitrogen  and  oxygen — air  will 

answer  every  purpose — the 
two  combine,  producing  the 
red  gas  nitrogen  peroxide, 
N02  (Fig.  29).  Under 
somewhat  similar  condi- 
tions— continual  supply  of 
energy  in  the  form  of  elec- 
tricity— nitrogen  can  even 
be  made  to  burn  in  air  with 
a  flame  which  is  hot  enough 
to  melt  a  platinum  wire. 

FIG.  S^mbin^ion  of  nitrogen  and  When  the  air  is  ^tensely 
oxygen  under  the  influence  of  a  stream  heated  in  any  Way — f  Or  in- 
o.  efcctric  sp.rks.  ^^  ^  ^^  m&g_ 

nesium  in  it — the  nitrogen  and  oxygen  combine  to  a  very 
perceptible  extent,  and  the  same  red  gas  is  produced. 

137.  Relation  of  nitrogen  to  life. — Nitrogen  is  an  im- 
portant element  in  organic  nature.  All  forms  of  life,  ani- 
mal or  vegetable,  contain  a  large  proportion  of  it,  and  the 
same  may  be  said  of  carbon,  hydrogen,  and  oxygen.  These 
four  are  preeminently  the  organic  elements.  However, 
many  other  elements  are  essential  constituents  in  living 
matter,  but  are  present  in  smaller  quantities. 

When  animal  matter  decays,  the  nitrogen  is  mostly  converted 
into  ammonia,  NH8.  About  one  tenth  of  the  nitrogen  escapes  in 
the  free  state.  The  ammonia  is  washed  into  the  soil  by  rains,  ab- 
sorbed by  plants,  and  converted  into  complex  nitrogenous  compounds 
which  serve  as  food  for  animals.  Some  plants,  like  peas,  beans,  and 
clover,  absorb  the  free  nitrogen  of  the  air  and  convert  it  into  complex 
compounds.  Such  plants  have  little  nodules  upon  their  roots,  com- 
posed of  certain  species  of  bacteria. 


CHAPTEE   XIV 

COMPOUNDS  OF  NITROGEN  AND  HYDROGEN 

138.  Ammonia,  NH3. — This  gas  is  a  product  of  the  de- 
cay of  organic  matter,  and  in  this  way  gets  into  the  at- 
mosphere, which  usually  contains  a  trace  of  it.     It  finds 
its  way  also  into  natural  waters,  and  the  presence  of  am- 
monia indicates  that  a  water  is  unfit  to  drink,  for  it  is 
evidence  of  recent  contamination  with  organic  matter,  prob- 
ably sewage. 

139.  Combination  of  nitrogen  and  hydrogen. — When  a 
mixture  of  nitrogen  and  hydrogen  is  confined  in  a  tube  over 
mercury  and  a  series  of  electric  sparks  passed  through  it, 
the  level  of  the  mercury  rises  and  partial  union  of  the  two 
to  form  ammonia  takes  place.    The  apparatus  used  is  that 
shown  in  Fig.  4. 

140.  Chemical  equilibrium, — The  combination  is  never  com- 
plete.    On  the  contrary,  if  we  confine  ammonia  gas  over  mercury 
and  treat  it  with  electric  sparks  in  the  same  way,  the  level  of  the 
mercury  falls  and  a  partial  separation  into  nitrogen  and  hydrogen 
occurs.     It  makes  no  difference  whether  we  start  with  ammonia  or 
with  the  mixture  of  nitrogen  and  hydrogen,  the  final  state  of  things 
will  be  the  same.     In  both  cases  the  gas  in  the  tube  will  contain 
the  same  percentages  of  hydrogen,  nitrogen,  and  ammonia  when  the 
process  is  over.     We  may  consider  that  two  processes  occur  together 
in  both  experiments — the  union  of  the  nitrogen  and  hydrogen  to 
ammonia,  and  the  separation  of  the  ammonia  into  its  constituents. 
Whether  there  is  observed,  on  the  whole,  separation  or  union,  will 
depend  on  whether  the  reaction, 

N  +  3H  =  NH3,  or 
NH3  =  N  +  3H, 

105 


106 


ELEMENTARY   CHEMISTRY 


occurs  more  rapidly.  When  the  two  have  the  same  speed — that  is, 
when  just  as  much  ammonia  is  decomposed  in  a  second  as  is  pro- 
duced in  the  same  time  there  is  equilibrium — the  process  is  sta- 
tionary. Many  chemical  changes  behave  in  this  way. 

141.  Source  of  the  ammonia  of  commerce. — Bituminous 
coal  always  contains  a  little  nitrogen,  and  when  it  is  distilled 
for  the  production  of  illuminating  gas,  the  nitrogen  unites 
with  some  of  the  hydrogen  which  such  coal  always  contains, 
producing  ammonia.  This  is  finally  obtained  in  a  very  im- 
pure form  in  solution  in  water,  and  this  liquid,  called  gas 

liquor,  is  the  source  of  near- 
ly all  the  ammonia  of  com- 
merce. 

142.  Preparation  of  am- 
monia gas  in  the  labora- 
tory.— The  most  convenient 
way  of  obtaining  ammonia 
gas  in  the  laboratory  is  to 
heat  ordinary  "  ammonia 
water,"  which  is  the  aque- 
ous solution  of  the  gas. 
The  ammonia  water  is 
placed  in  a  flask  with  a  de- 
livery tube  and  heated 
gently  (Fig.  30).  The  gas 
can  not  be  collected  over  water,  for  it  is  very  soluble  in  that 
liquid:  For  this  reason,  so  long  as  water  was  the  only 
liquid  used  by  chemists  in  collecting  gases,  ammonia  re- 
mained undiscovered ;  and  when,  in  1774,  Priestley  used  mer- 
cury in  place  of  the  water,  its  discovery  was  one  of  the  first 
results  of  the  new  method.  On  account  of  its  great  weight, 
mercury  can  not  well  be  used  for  collecting  large  quantities, 
and  there  is  nothing  left  but  the  method  of  dry  displace- 
ment, which  we  have  already  employed  in  the  case  of  chlo- 
rine. But  ammonia  is  much  lighter  than  air.  Hence  we 


Flo.  30.— Preparation  of  ammonia  gas 
from  ammonia  water. 


COMPOUNDS  OF  NITROGEN  AND   HYDROGEN      107 

invert  the  jar  in  which  we  wish  to  collect  it,  and  pass  the 
tube  conveying  the  gas  up  to  the  bottom  of  the  jar.  The 
displaced  air  flows  out  around  the  mouth  of  the  jar. 

143.  Properties.  —  Ammonia  is  a  colorless  gas,  with  a 
sharp,  peculiar  odor.     It  is  not  poisonous,  but  its  action 
on  the  mucous  membrane  is  irritating,  and  the  prolonged 
breathing  of  air  containing  small  quantities  is  attended  with 
bad  results.     It  is  easily  converted  by  cold  or  pressure,  or 
by  both  together,  into  a  colorless  liquid.     This  liquid  pro- 
duces great  cold  when  it  evaporates,  and  is  largely  em- 
ployed for  cooling  and  for  the  production  of  artificial  ice, 
for  which  purpose  it  has  almost  completely  displaced  other 
substances  in  practical  work. 

144.  Combustion  of  ammonia  in  air  and  oxygen.  —  Am- 
monia will  not  burn  with  a  continuous  flame  in  the  air. 
If  a  tube  through  which  the  gas  is  issuing  is  held  near  a 
lighted  burner,  the  ammonia  burns  around  the  gas  flame  as 
a  pale  yellow  mantle,  but  is  extinguished  when  the  Bunsen 
flame  is  removed.     In  oxygen,  on  the  contrary,  it  burns 
steadily  with  a  peculiar  yellow  flame.    The  hydrogen  burns 
to  water  and  the  nitrogen  is  liberated: 


30  =  N2  +  3H20. 

Ammonia  is  one  of  the  very  soluble  gases.  One  vol- 
ume of  water  absorbs  at  0°,  and  under  the  pressure  of 
one  atmosphere  over  1,100  volumes  of  it.  Ammonia 
water,  or  "hartshorn,"  is  an  aqueous  solution  of  the 
gas.  This  solution  is  colorless  and  has  the  odor  of  the 
gas.  When  swallowed,  it  acts  as  a  powerful  irritant  poi- 
son. It  is  largely  used  in  the  laboratory  and  in  the 
household. 

145.  Ammonium  compounds.  —  A  cylinder  is  filled  with 
ammonia  and  another  similar  one  with  hydrochloric  acid. 
Both  are  covered  with  glass  plates.  The  cylinders  are 
brought  mouth  to  mouth  and  the  plates  removed,  so  that 


108  ELEMENTARY  CHEMISTRY 

the  gases  mix.    At  once  a  dense  white  smoke  fills  the  in- 
terior of  both  vessels,  and  in  a  little  while  deposits  as  a 

white    film    on    the    glass    (Fig. 

31).      This   is   ammonium   chlo- 

ride, NH4C1,  produced  by  combi- 

nation : 

NH3  +  HC1  =  NH4C1. 

It  is  worth    noting   that,  if  both 
gases  are  completely  free  from  water, 
they  will  remain  clear  when  mixed,  no 
ammonium  chloride  being  produced. 
FIG.  31.—  Combination  of  ammonia    The  presence   of  the  smallest  trace  of 
gas  with  hydrochloric-acid  gas.       water.vapor  giveg  rise  at  Qnce  to  the 

combination  in  which  it  takes  no  part.     This  is  a  remarkable  case  of 
catalytic  action. 

Ammonia  combines  with  other  acids,  producing  com- 
pounds in  which  the  hydrogen  of  the  acid  is  replaced  by  the 
group  NH4,  and  which  are  called  ammonium  salts  — 


8  as  NH4N03, 

nitric  acid  ammonium  nitrate 


H2S04  =  (NH4)2S04, 

sulphuric  acid  ammonium  sulphate 

and  so  on.  These  salts  are  white,  unless  the  acid  combined 
with  the  ammonia  is  colored,  and  are  soluble  in  water. 
When  heated  many  of  them  are  completely  converted  into 
vapor,  leaving  no  residue,  and  this  vapor  is  usually  not  that 
of  the  unchanged  salt,  but  a  mixture  of  ammonia  and  the 
acid.  Thus,  when  ammonium  chloride  is  heated,  it  disap- 
pears completely,  and  the  vapor  is  a  mixture  of  ammonia 
and  hydrochloric  acid  in  equal  volumes  : 

HC1. 


If  the  ammonium  chloride  is  absolutely  dry  it  is  converted  into 
vapor  without  this  separation. 


• 

COMPOUNDS  OF  NITROGEN  AND  HYDROGEN      109 

Every  acid  yields  its  corresponding  ammonium  salt,  but 
the  three  whose  formulas  have  been  given  are  the  most 
important.  Ammonium  chloride  has  a  healing  action  on 
inflamed  mucous  membranes,  and  is  a  frequent  ingredient 
in  cough  mixtures  and  lozenges.  Ammonium  sulphate  is 
placed  on  the  soil  as  a  source  of  nitrogen  to  crops,  and 
is  used  extensively  in  the  manufacture  of  alum.  Ammoni- 
um nitrate  is  employed  in  the  manufacture  of  nitrous  oxide 
gas,  so  much  used  by  dentists.  All  three  are  white  and 
crystalline,  and  readily  soluble  in  water. 

146.  Radicals. — We  have  learned  to  regard  the  formula 
NaCl  as  meaning  a  compound  containing  23  parts  of  sodium 
and  35.5  parts  of  chlorine  by  weight.  We  can  now  think  of 
it  also  as  meaning  a  molecule  containing  an  atom  of  chlorine 
weighing  35.5  and  one  of  sodium  weighing  23.  Now,  the 
compound  NH4C1  is  also  a  chloride,  but  instead  of  contain- 
ing a  single  atom  like  Na  united  with  the  Cl,  it  contains 
a  group  of  atoms  consisting  of  one  nitrogen  and  four  hydro- 
gen atoms,  a  group  which  we  indicate  therefore  by  the  for- 
mula NH4.  This  group,  which  is  able  to  fill  the  place  of  the 
sodium,  and  to  form  a  chloride  in  which  it  exists  in  union 
with  Cl,  just  like  an  element,  is  called  ammonium,  and  the 
term  radical  is  applied  as  a  class  name  to  all  such  groups. 
Thus,  ammonium,  NH4,  is  a  radical,  and  all  the  compounds 
in  which  we  assume  it  to  exist  are  called  ammonium  com- 
pounds, just  as  compounds  containing  Na  are  called  sodium 
compounds.  Many  such  compounds  are  known,  e.  g. : 

Ammonium  Sodium 

compounds.  compounds. 

NH4C1  (ammonium  chloride) NaCl. 

NH4OH  (ammonium  hydroxide) NaOH. 

NH4Br  (ammonium  bromide) NaBr. 

(NH4)aCO8  (ammonium  carbonate) Na9CO>. 

NH4NO,  (ammonium  nitrate) NaNO,. 

Notice,  that  in  this  list — and  it  might  be  much  extended 
— the  likeness  between  the  ammonium  and  sodium  com- 


110  ELEMENTARY  CHEMISTRY 

pounds  is  complete,  provided  we  assume  that  NH4,  one 
atom  of  nitrogen  and  four  of  hydrogen,  plays  the  same  role 
in  the  molecule  as  Na,  one  atom  of  sodium,  in  the  molecule 
of  the  corresponding  sodium  compound. 

Many  other  radicals  are  known.  We  have  noticed  that 
sodium  hydroxide,  NaOH,  and  ammonium  hydroxide, 
NH4OH,  both  contain  the  complex  OH.  All  hydroxides 
contain  it,  for  example: 

Lithium  hydroxide,  LiOH; 
Potassium  hydroxide,  KOH; 
Calcium  hydroxide,  Ca(OH)2. 

OH  is  therefore  a  radical,  and  is  called  hydroxyl. 

147.  Ammonium  hydroxide,  NH4OH,  is  produced  in 
small  quantities,  when  ammonia,  N"H3,  dissolves  in  water: 

H20  =  NH4OH. 


It  is  very  unstable,  and  has  never  been  obtained  in  the 
solid  state. 

In  addition  to  ammonia,  two  other  compounds  of  hydrogen  and 
nitrogen  have  been  prepared. 

Hydrazine,  NalU,  is  a  colorless  liquid  with  an  ammoniacal  odor. 
When  heated  it  boils,  passing  into  a  colorless  vapor.  Like  ammo- 
nia, it  combines  with  acids. 

Hydrazoic  acid,  N3H,  is  a  colorless  gas  with  a  penetrating,  intol- 
erable odor.  It  is  very  soluble  in  water,  and  the  solution  is  strongly 
acid,  dissolving  many  of  the  metals  —  zinc  and  iron,  for  instance  — 
with  evolution  of  hydrogen. 


CHAPTER    XV 

COMPOUNDS   CONTAINING   NITROGEN  AND    OXYGEN 

14*8.  Four  oxides  of  nitrogen  are  known : 

Nitrous  oxide,  N20; 

Nitric  oxide,  NO; 

Nitrogen  peroxide,  N02(N204)  ; 

Nitrogen  pentoxide,  N205. 

Another  compound,  N203 — nitrogen  trioxide — is  often 
described,  but  its  existence  is  doubtful. 

149.  Nitrous  oxide,  N20,  is  made  by  gently  heating  am- 
monium nitrate  in  a  glass  flask : 

NH4N03  =  N20  +  2H20. 

The  gas  can  be  collected  over  water,  though  some  of  it  will 
dissolve  and  be  lost. 

Nitrous  oxide  is  a  colorless  gas,  with  a  slight  pleasant 
odor  and  a  faint  sweet  taste.  Water  dissolves  a  little  more 
than  its  own  volume  under  ordinary  temperature  and  pres- 
sure. The  gas  is  readily  converted  by  cold  or  pressure  into 
a  colorless  liquid.  This  liquid  is  made  in  large  quantities, 
and  is  sold  in  strong  steel  cylinders  for  the  use  of  dentists. 
When  nitrous  oxide  is  inhaled  brief  insensibility  results,  and 
this  fact  is  utilized  in  the  extraction  of  teeth. 

Nitrous  oxide  is  decomposed  into  its  elements  by  heat. 
This  fact  is  the  key  to  its  peculiar  behavior  toward  com- 
bustible substances.  An  instructive  experiment  in  this  con- 
nection is  to  place  a  piece  of  phosphorus  in  nitrous'  oxide 

111 


112  ELEMENTARY  CHEMISTRY 

and  touch  it  with  a  hot  metallic  rod.  The  phosphorus  melts 
and  vaporizes  at  the  point  touched,  but  does  not  catch  fire 
because  there  is  no  free  oxygen  present.  But  if  the  phos- 
phorus is  placed  in  air  and  its  combustion  started  and  then 
plunged  into  the  gas,  it  burns  brilliantly,  because  the  heat 
of  the  flame  separates  the  nitrous  oxide  into  its  elements, 
and  the  free  oxygen  supports  the  combustion.  Sulphur 
feebly  burning  is  extinguished  when  placed  in  nitrous  oxide, 
but  if  the  sulphur  is  burning  vigorously  its  combustion  con- 
tinues. 

150.  Nitric  oxide,  NO,  is  made  by  means  of  the  same 
apparatus  which  we  employed  in  making  hydrogen  (Fig. 
10).  Copper  turnings  covered  with  water  are  placed  in  the 
gas  bottle  and  nitric  acid  gradually  added  through  the  fun- 
nel tube.  The  gas  in  the  bottle  is  reddish  brown  at  the 
start  because  the  NO  which  is  liberated  at  first  combines 
with  the  oxygen  of  the  air  and  produces  the  strongly  col- 
ored gas  N02.  Soon  this  is  swept  out  and  the  gas  becomes 
colorless.  This  is  nitric  oxide.  It  can  be  collected  over 
water,  in  which  it  is  almost  insoluble.  The  equation  is : 

3Cu  +  8HN03  =  3Cu(N03)2  +  4H20  +  2NO. 

Nitric  oxide  is  a  colorless  gas  which  can  be  converted  by 
pressure  and  cold  into  a  colorless  liquid.  Its  odor  is  un- 
known because,  as  soon  as  it  comes  in  contact  with  air,  it 
combines  with  more  oxygen  and  produces  the  strongly 
smelling  gas  N02.  For  the  same  reason  we  know  nothing 
regarding  its  action  on  the  body.  It  is  impossible  to  inhale 
it,  for  it  would  come  into  contact  with  free  oxygen  in  the 
lungs,  and  there  produce  N02,  which  is  violently  poisonous. 
In  order  to  avoid  this,  Sir  Humphry  Davy  filled  his  lungs 
first  with  nitrous  oxide  and  then  inhaled  nitric  oxide.  The 
result  was  a  burning  sensation  in  the  throat  so  intense  that 
the  experiment  was  discontinued. 

So  far  as  supporting  combustion  is  concerned,  the  be- 


NITROGEN  AND  OXYGEN  113 

havior  of  nitric  oxide  is  somewhat  the  same  as  that  of  nitrous 
oxide,  but  a  higher  temperature  is  required  to  separate  it 
into  its  constituents,  and  it  extinguishes  some  combustibles, 
like  sulphur  and  a  candle,  which  will  continue  to  burn  in 
nitrous  oxide.  Yet  a  substance  which  will  burn  in  it  —  phos- 
phorus or  magnesium,  for  instance  —  burns  more  brilliantly 
than  in  nitrous  oxide  because  it  is  more  liberally  supplied 
with  oxygen.  The  equations  — 

2N0  =  2N        0 


show  that  in  the  decomposition  of  nitrous  oxide  the  mixture 
produced  contains  one-third  of  its  volume  of  oxygen;  in  the 
case  of  nitric  oxide,  one-half. 

151.  Nitrogen  peroxide,  N02(N204).  —  Some  lead  nitrate 
is  placed  in  a  retort  and  heated  gently.  The  gases  given 
off  are  passed  into  a  U-shaped  tube  surrounded  by  a  mix- 
ture of  ice  and  salt  (Fig.  32).  The  lead  nitrate  is  decom- 
posed according  to  the  following 
equation  : 

Pb(N03)2  =  PbO  +  2N02  +  0. 

The  nitrogen  peroxide  condenses 
in  the  U-tube,  the  oxygen  pass- 
ing on  unchanged. 

Nitrogen  peroxide  is,  at  low 
temperatures,  a  white  solid  which 

melts    to    a    COlorleSS    liquid.       If     FIG.  32.—  Preparation  of  nitrogen 

the    temperature    is    allowed    to  peroxide'  NO" 

rise  the  liquid  becomes  yellow,  and  is  finally  converted  into 
a  gas  which  at  —  10°  is  faint  yellow,  nearly  colorless.  If 
the  gas  be  heated  the  tint  deepens,  until  at  140°  it  possesses 
a  reddish-black  color  so  intense  that  a  layer  2  centimetres 
thick  is  opaque. 

This  variation  in  properties  with  the  temperature  is  surprising, 
but  its  cause  becomes  clear  when  the  density  of  the  gas  is  deter- 


114  ELEMENTARY  CHEMISTRY 

mined  at  different  temperatures.  When  this  is  done  it  is  found  that 
the  molecular  weight  of  the  colorless  gas  corresponds  to  the  formula 
NaO4,  while  the  deep  reddish-black  gas  obtained  at  140°  has  a 
molecular  weight  only  half  as  great  and  possesses  the  formula  NOa. 
Nitrogen  peroxide,  then,  exists  in  two  conditions :  as  NaO4,  which 
is  colorless,  has  the  molecular  weight  92,  and  the  density  46  referred 
to  hydrogen ;  and  as  NOa.  which  is  reddish  black,  has  the  molecu- 
lar weight  46,  and  the  hydrogen  density  23.  When  the  substance 
is  heated,  the  reaction — 

.  N3O4  =  2NOa— 

occurs,  and  the  color  becomes  deeper  as  the  temperature  rises, 
because  the  quantity  of  NOa  in  the  gas  increases.  Finally,  at  140°, 
the  change  is  complete,  and  then  heating  to  a  higher  temperature 
does  not  make  the  color  more  intense.  On  cooling,  the  change  is 
reversed.  This  is  an  interesting  case  of  dissociation,  or  ' '  gradual 
decomposition,  increasing  with  the  temperature  and  reversed  by 
cooling."  It  is  important  to  understand  clearly  the  distinction 
between  NaO4  and  NO-j.1  Both  have  exactly  the  same  composition. 
If  we  think  of  matter  in  mass,  then  the  difference  is  that  N2O4  has 
a  density  twice  as  great  as  that  of  NO2.  If  we  think  of  molecules, 
then  the  molecule  NaO4  is  twice  as  heavy.  We  express  this  by  say- 
ing that  NaO4  is  a  polymer  of  NOa.  It  will  be  seen  that  we  have 
already  studied  a  very  similar  state  of  things  in  the  relation  between 
ozone  and  oxygen. 

152.  Nitrogen  pentoxide,  N205,  can  be  obtained  by  ab- 
stracting the  elements  of  water  from  nitric  acid,  HN03: 

2HN03-HoO  =  N205. 

—       '  "L-  ^£~  ~/~    *!    *i<-  £/ 

In  preparing  nitrogen  peritoxide,  pure  nitric  acid  is 
mixed  with  phosphorus  pentoxide,  P205,  and  the  mixture 
distilled.  Nitrogen  pentoxide  is  a  colorless  crystalline  sub- 
stance which  is  very  unstable,  and  explodes  spontaneously 
on  being  preserved.  When  thrown  into  water  the  liquid 
becomes  hot,  and  is  found  to  contain  nitric  acid : 

N205  +  H20  =  2HN03. 
1  This  will  be  better  understood  after  reading  Chapter  XVI. 


NITROGEN   AND  OXYGEN 


115 


153.  Nitric  acid,  HN03,  is  obtained  by  distilling  a  mix- 
ture of  sodium  nitrate,  N"aN"03,  with  sulphuric  acid.  In  the 
laboratory  the  reaction  can  be  carried  out  in  the  apparatus 
shown  in  Fig.  33.  The  vapor  of  the  acid  liberated  in  the 
retort  is  led  into  a  glass  vessel  around  which  cold  water  cir- 
culates. Here  it  condenses  to  a  liquid. 


FIG.  33.— Preparation  of  nitric  acid. 

On  the  large  scale,  cylinders  of  cast  iron  are  employed 
to  contain  the  mixture  of  sodium  nitrate  and  sulphuric  acid, 
and  the  condensation  of  the  nitric  acid  is  effected  in  stone- 
ware bottles.  The  equation  is : 

2NaN03  +  H2S04  =  Na2S04  +  2HN"03. 

154.  Properties. — Pure  nitric  acid  is  a  colorless,  fuming 
liquid,  about  one  and  one-half  time  as  heavy  as  water.  Its 
vapor  has  an  acrid  odor  and  is  poisonous,  and  the  liquid 
attacks  the  skin  and  produces  all  the  effects  of  a  violent 
irritant  poison  when  swallowed.  It  attacks  many  metals, 


116  ELEMENTARY   CHEMISTRY 

converting  some,  like  copper  and  silver,  into  their  nitrates 
(e.  g.,  silver  nitrate,  AgN03),  others,  like  tin,  into  their 
oxides  (e.g.,  Sn02).  Still  others,  like  gold  and  platinum, 
are  not  affected  by  it.  Hydrogen  is  never  liberated  by  the 
action  of  nitric  acid  upon  a  metal.  Usually  the  gas  given 
off  is  NO  or  N02.  With  cold  dilute  acid,  N20  is  obtained. 

155.  Uses. — Mtric  acid  is  a  commercial  product  of  some 
importance.    Considerable  quantities  of  it  are  employed  in 
the  manufacture  of  nitroglycerin  and  guncotton,  and  the 
use  of  the  acid  for  this  purpose  has  increased  of  late  on 
account  of  the  production  of  smokeless  gunpowders,  most 
of  which  are  mixtures  containing  guncotton  or  guncotton 
and  nitroglycerin.     It  is  also  employed  in  etching  metallic 
surfaces,  and  to  some  extent  in  the  separation  of  gold  and 
silver.     From  an  alloy  of  these  two  metals  it  dissolves  the 
silver  to  nitrate  (AgN03),  leaving  the  gold  unaffected.    At 
present,  sulphuric  acid  is  replacing  it  in  this  operation. 

156.  The  nitrates. — The  salts  of  nitric  acid  are  called 
the  nitrates,  and,  like  the  salts  of  all  acids,  each  consists  of 
the  acid  with  its  hydrogen  replaced  by  some  metal,  e.  g., 

Nitric  acid 

HN03.  KN03,  potassium  nitrate; 

NaN03,  sodium  nitrate; 
Cu(N03)2,  copper  nitrate; 
Bi(N03)3,  bismuth  nitrate. 

The  nitrates  are  all  soluble  in  water  and  are  all  decom- 
posed by  heat,  leaving  a  residue,  which  usually  consists  of 
the  oxide  of  the  metal  whose  nitrate  was  heated.  If  this 
oxide  is  itself  decomposed  by  heat,  the  metal  is  obtained  in- 
stead. Thus,  silver  nitrate  first  melts,  then  gives  off  N02 
and  oxygen,  and  leaves  a  residue  of  silver. 

157.  Potassium  nitrate,,  KN03,  commonly  called  salt- 
peter, occurs  as  an  incrustation  on  the  soil  in  hot  countries. 
It  consists  of  white  crystals,  very  soluble  in  water,  which 
melt  easily  when  heated  to  a  colorless  liquid.     This  salt 


NITROGEN  AND   OXYGEN  117 

was  formerly  made  in  large  quantities  for  the  manufacture 
of  black  gunpowder.  This  is  a  mixture  of  about  the  fol- 
lowing composition : 

Potassium  nitrate 75  per  cent ; 

Charcoal 15        " 

Sulphur    10        " 

Black  gunpowder  is  still  largely  employed  for  saluting 
purposes  and  for  hunting.  In  warfare  it  is  obsolete. 

158.  Sodium  nitrate,  NaN03,  is  by  far  the  most  im- 
portant salt  of  nitric  acid  commercially.     In  the  rainless 
region  in  the  northern  part  of  Chile,  near  the  junction  with 
Peru  and  Bolivia,  immense  deposits  of  this  salt  occur,  and 
about  1,500,000  tons  of  it  are  exported  yearly.    About  three- 
fourths  of  this  are  employed  as  a  fertilizer,  in  order  to 
furnish  nitrogen  to  crops;  the  remainder  goes  into   the 
chemical  industries,  being  used  chiefly  in  the  manufacture 
of  nitric  acid  and  of  potassium  nitrate. 

Sodium  nitrate  is  white,  and  is  more  soluble  in  water 
than  potassium  nitrate.  It  deliquesces  in  moist  air,  which 
is  not  the  case  with  the  potassium  salt. 

159.  Nitrous    acid,   HN02. — When    sodium    nitrate    is 
carefully  heated,  one-third  of  its  oxygen  escapes  and  the 
residue  has  the  composition  NaN02.    It  is  called  sodium  ni- 
trite.   Potassium  nitrate  behaves  in  the  same  way,  leaving  a 
residue  of  potassium  nitrite,  KN02.     Many  other  nitrites 
are  known,  most  of  which,  like  potassium  and  sodium  ni- 
trites, are  colorless  or  faint  yellow,  and  freely  soluble  in 
water.    Silver  nitrite  is  only  slightly  soluble,  arid  separates 
as  a  white  precipitate  when  solutions  of  potassium  nitrite 
and  silver  nitrate  are  mixed: 

KN02  +  AgN03  =  AgN02  +  KN"03. 

These  compounds  are  regarded  as  salts  of  nitrous  acid, 
HN02,  but  the  acid  itself  has  never  been  prepared.  It  is 


118  ELEMENTARY  CHEMISTRY 

only  known  dissolved  in  water,  and  when  we  attempt  to 
obtain  it  by  evaporating  the  liquid  it  decomposes : 

2HN02  =  H20  +  NO  +  N02. 

160.  Compounds  of  nitrogen  with  chlorine,  bromine,  and 
iodine. — When  chlorine  is  passed  into  a  solution  of  ammo- 
nium chloride,  nitrogen  chloride,  NC13,  is  produced: 

NH4C1  +  3C12  =  NCI,  +  4HC1. 

Nitrogen  chloride  is  a  yellow  oil  with  a  penetrating  odor. 
Its  vapor  attacks  the  mucous  membranes,  and  by  prolonged 
inhalation  produces  permanent  inflammation.  The  liquid 
explodes  violently  on  being  heated  and  also  on  contact  with 
many  substances;  for  instance,  with  rubber  or  oil  of  tur- 
pentine. Sometimes  it  seems  to  explode  spontaneously. 
The  explosion  is  simply  a  sudden  separation  into  nitrogen 
and  chlorine,  a  process  which  is  accompanied  by  great  ex- 
pansion and  liberation  of  heat. 

Nitrogen  bromide,  NBr8,  is  very  imperfectly  investigated.  It  is 
said  to  be  a  red  oil  which  explodes  like  the  chloride. 

The  analogous  compound  with  iodine  would  be  NI8.  This  is 
unknown,  but  a  compound,  N2H,IS,  has  been  studied,  and  this  we 
can  regard  as  consisting  of  a  molecule  of  NI8  combined  with  a 
molecule  of  ammonia;  thus,  NH3,NI8.  It  forms  dark  crystals  with 
a  coppery,  metallic  luster.  It  is  extremely  explosive. 


CHAPTER   XVI 

ATOMIC  AND  MOLECULAR   WEIGHTS- AVOGADR&S  RULE 

161.  The  molecular  weight  of  HC1. — When  we  analyze 
hydrochloric  acid  we  find  that  it  contains  for  1  part  of 
hydrogen  35.5  parts  of  chlorine.  If,  then,  we  assume  that  it 
contains  one  atomic  weight  of  each  element,  it  follows  at 
once  that  the  atomic  weight  of  chlorine  is  35.5.  It  would 
lead  us  too  far  to  discuss  here  the  reasons  for  accepting  this 
as  the  real  atomic  weight  of  this  element.  We  will  content 
ourselves  with  the  statement  that,  by  assigning  the  value 
35.5  to  the  atomic  weight  of  chlorine,  we  can  write  simple 
formulas  for  all  the  chlorine  compounds.  We  are  never  com- 
pelled to  use  fractional  parts  of  a  symbol  in  writing  these 
formulas. 

HC1  is,  then,  the  formula  of  hydrochloric  acid,  and  when 
we  add  together  the  parts  by  weight  indicated  in  the  formu- 
la, we  obtain : 

H=   1. 

Cl  =  35.5 


HC1  =  36.5 

This  figure,  36.5,  we  call  the  molecular  weight  of  the  sub- 
stance. From  the  standpoint  of  the  atomic  theory,  HC1 
means  an  atom  of  hydrogen  in  union  with  an  atom  of  chlo- 
rine, forming  a  molecule  of  hydrochloric  acid,  and  this  mole- 
cule must  weigh  36.5  times  as  much  as  the  hydrogen  atom. 
We  do  not  know  what  the  atom  of  hydrogen  weighs,  but  if 
we  assume  it  to  weigh  one  ten-millionth  of  a  milligram,  then 

119 


120  ELEMENTARY  CHEMISTRY 

the  molecule  HC1  would  weigh  36.5  ten-millionths  of  a 
milligram. 

Returning  to  grams,  the  formula  HC1  means  36.5  grams 
of  a  substance  which  consists  of  1  gram  of  hydrogen  and 
35.5  grams  of  chlorine.  The  volume  which  these  36.5  grams 
of  hydrochloric-acid  gas  will  occupy  depends  upon  the  tem- 
perature and  the  pressure;  at  the  standard  temperature,  0°, 
and  the  standard  pressure,  760  millimeters  of  mercury,  the 
volume  is  22.4  liters. 

Now,  the  volume  occupied  by  the  molecular  weights  in 
grams  of  all  gases  and  vapors  is  the  same,  at  the  same  tem- 
perature and  pressure;  and  this  volume  is  224  liters  if  the 
temperature  is  0°  and  the  pressure  760  millimeters.  For 
example  : 


Name  of  gas.  Fonnuta. 

Hydrochloric  acid  ................  HC1  36.5  22.4  liters. 

Ammonia  .......................  NH,  17  22.4      " 

Nitrous  oxide  ....................  NaO  44  22.4     '" 

Nitric  oxide  .................  ....  NO  30  22.4      " 

162.  Calculations  based  on  the  rule  stated  above.  —  This 
fact,  that  the  volume  of  the  molecular  weight,  taken  in 
grams,  is  the  same  for  all  gases  and  vapors  —  22.4  liters  at 
0°  and  760  millimeters  —  gives  us  a  simple  method  of  cal- 
culating the  weight  of  one  liter  of  any  gas.  We  have  only 
to  divide  the  molecular  weight  by  22.4.  Thus,  the  weight  of 

1  liter  of  nitrous  oxide,  N20,  is  —  =  1.9642  gram.    The 

22.4 

same  fact  enables  us  to  determine  at  once  from  the  formula 
the  density  of  any  gas  or  vapor  referred  to  hydrogen  as 
unity.  The  formula  of  hydrogen  is  H2.  The  molecular 
weight  is  therefore  2,  and  it  requires  2  grams  of  the  gas 
to  fill  the  same  volume  as  the  molecular  weights  of  other 
gases.  But  the  density  of  any  gas  referred  to  hydrogen  is 
simply  the  weight  of  any  volume  of  it  divided  by  the  weight 
of  the  same  volume  of  hydrogen.  The  density  is  therefore 


AVOGADRO'S  RULE 

equal  to  the  molecular  weight  divided  by  the  2  grams  of 
hydrogen  which  fill  the  same  bulk.  If  D  —  density  and 
M  =r  molecular  weight,  we  have  therefore — 

M 

D=  —  - 

2 

Thus,  the  molecular  weight  of  nitrous  oxide  is  44  and  the 

44 

density  is  -—  =  22.     Any   volume   of  the  gas  weighs   22 
Z 

times  as  much  as  the  same  volume  of  hydrogen,  both  being 
at  the  same  temperature  and  pressure. 

163.  The  molecular  weight  can  be  calculated  by  multi- 
plying the  density  by  2. — Of  course,  if  we  determine  the 
density  of  a  gas  or  vapor  by  experiment,  we  can  calculate 
the  molecular  weight  by  multiplying  the  density  by  2,  and 
this  is  a  most  important  method  of  ascertaining  the  molec- 
ular weights  of  new  substances  whose  formulas  have  not 
jet  been  established.     By  weighing  a  measured  volume  of 
pure  nitrous  oxide  we  could  easily  show  that  its  density,  re- 
ferred to  hydrogen,  is  about  22,  and  from  this  it  follows 
that  the  molecular  weight  is  22  X  2,  or  44. 

164.  Avogadro's  rule. — We  have  just  stated  that  the 
molecular  weights  of  all  gases  and  vapors  occupy  the  same 
volume,  temperature  and  pressure  being  the  same.    Let  us 
now  translate  this  statement  into  the  language  of  atoms 
and  molecules.    In  the  first  place,  it  is  clear  that  the  molecu- 
lar weights,  taken  in  grams,  of  all  substances,  must  contain 
the  same  number  of  molecules. 

In  order  to  understand  this  statement,  think  of  any  substance 
which  consists  of  individual  grains  all  of  which  are  alike — of  shot, 
for  instance.  Let  us  consider  that  we  have  several  different  kinds 
o'f  shot,  and  that  we  are  required  to  choose  quantities  by  weight  of 
the  different  kinds  which  will  contain  the  same  number  of  pellets. 
We  will  take  one  decigram  (.1  gram)  as  our  unit  of  weight,  and 
suppose  that  each  grain  of  the  first  kind  of  shot  weighs  2  decigrams 
(.2  gram).  Then  2  kilos  will  contain  ^f^,  or  10,000  pellets.  Each 


122  ELEMENTARY  CHEMISTRY 

grain  of  a  second  sort  weighs  10  decigrams  or  1  gram.  How  much 
of  this  variety  will  be  required  to  contain  10,000  pellets  ?  Clearly 
10  kilos,  for  each  kilo  will  contain  1,000.  Finally,  if  we  imagine 
a  third  variety,  consisting  of  bullets  each  of  which  weighs  100 
decigrams,  or  10  grams,  it  is  plain  that  in  order  to  obtain  10,000 
bullets  we  must  weigh  off  100  kilos,  for  each  kilo  will  contain  ^1-, 
or  100. 

What  is  true  of  the  visible  shot  is  equally  true  of  the  invisible 
molecules. 

17  grams  of  ammonia,  NH3, 
44       "        "    nitrous  oxide,  N20, 
30       "        "    nitric  oxide,  NO, 

2       "        "    hydrogen,  H2, 

must  all  contain  the  same  number  of  molecules.  But  we 
have  seen  that  these  quantities  occupy  equal  volumes. 
Hence,  in  these  equal  volumes  equal  numbers  of  molecules 
are  contained.  In  general,  the  argument  is  this : 

1.  The  molecular  weights  in  grams  of  all  substances 
contain  equal  numbers  of  molecules. 

2.  The  molecular  weights  in  grams  of  all  gases  and  va- 
pors occupy  equal  volumes  at  the  same  temperature  and 
pressure.    Therefore — 

3.  In  equal  volumes  of  all  gases  and  vapors  at  the  same 
temperature  and  pressure,  equal  numbers  of  molecules  are 
contained.     This  statement  is  commonly  called  Avogadro's 
rule. 

The  density  of  a  gas  is  the  weight  of  any  volume  of  it 
divided  by  the  weight  of  the  same  volume  of  hydrogen.  It 
is  therefore  the  weight  of  a  certain  number  of  molecules 
of  the  gas  divided  by  the  weight  of  the  same  number  of 
molecules  of  hydrogen.  But  this  is  the  same  thing  as  the 
weight  of  one  molecule  of  the  gas  divided  by  the  weight  of 
one  molecule  of  hydrogen.  Thus  the  density  of  nitrous 
oxide  is — 

Weight  of  one  molecule,  N20      44 
Weight  of  one  molecule,  H2  ~~  2 


AVOGADRO'S  RULE  123 

But,  since  all  our  atomic  weights  are  in  terms  of  the  weight 
of  the  hydrogen  atom  as  1,  the  molecular  weight  of  nitrous 
oxide  is — 

Weight  of  one  molecule,  N20       44 

-        --  -  _.  _..  — — -    •  A  A 

Weight  of  one  atom,  H  1 

Again,  we  see  that  the  density  is  half  the  molecular  weight. 
In  determining  the  molecular  weights  our  unit  is  the  hy- 
drogen atom,  H  =  1 ;  in  determining  the  densities,  it  is  the 
hydrogen  molecule,  H2  =  2. 

165.  The  law  of  simple  volume  ratios. — We  have  no- 
ticed (p.  80)  that,  when  two  gases  combine,  there  is  always 
a  simple  relation  between  the  volumes  of  the  gases  entering 
into  the  change.    This  is  just  what  -the  atomic  theory  would 
lead  us  to  expect.     Chemical  combinations  take  place  be- 
tween molecules,  and  the  number  of  molecules  of  each  sub- 
stance entering  into  the  change  is  always  small.     Equal 
volumes  of  gases  contain  equal  numbers  of  molecules.    If 
one  molecule  of  one  gas  combines  with  one  molecule  of  the 
second,  then  one  volume  of  the  first  will  combine  with  one 
volume  of  the  second,  thus — 

H2  +  Cla  =  2HC1. 

On  the  other  hand,  in  the  production  of  water,  two 
molecules  of  hydrogen  react  with  one  of  oxygen ;  therefore, 
two  volumes  of  hydrogen  react  with  one  of  oxygen: 

2H2  +  02  =  2H20. 

In  other  words,  so  far  as  gases  and  vapors  are  con- 
cerned, the  formulas  in  the  equation,  which  represent  mole- 
cules, can  be  also  read  as  representing  volumes. 

166.  Doubled  formulas. — Finally,  we  must  inquire  why 
it  is  that  in  some  cases  we  express  the  composition  of  sub- 
stances by  formulas  which  are  multiples  of  the  simplest 
expressions  which  would  serve  that  purpose.     Thus,  for 
hydrogen  peroxide  we  write  H202,  for  hydrazine  N2H4, 


124  ELEMENTARY  CHEMISTRY 

when  the  formulas  OH  and  NH2  would  mean  exactly  the 
same  thing,  so  far  as  chemical  composition  goes. 

A  doubled  formula  means  usually  that  the  density  of 
the  gas  or  vapor  has  been  determined,  and  from  this,  by 
multiplying  by  2,  the  molecular  weight ;  and  that  the  latter 
has  been  found  to  be  twice  that  which  the  simplest  formula 
requires.  Thus,  there  is  a  gas  called  cyanogen  which  con- 
tains 12  parts  of  carbon  in  union  with  14  of  nitrogen.  Its 
composition  is  therefore  accurately  expressed  by  the  formu- 
la CN".  But  its  density  referred  to  hydrogen  is  26,  and 
therefore  the  molecular  weight  must  be  26  X  2,  or  52.  Ac- 
cordingly, we  must  write  for  the  gas  the  doubled  formula 
C2N2. 

With  hydrogen  peroxide  we  can  not  ascertain  the  molecu- 
lar weight  in  this  way,  because  when  we  attempt  to  convert 
it  into  vapor  to  obtain  the  density,  it  separates  into  oxygen 
and  water.  Many  other  substances  behave  in  the  same  way. 
But  within  the  last  few  years  methods  have  been  devised  by 
which  it  is  possible  to  determine  the  molecular  weights  of 
dissolved  substances.  Thus,  when  hydrogen  peroxide  is  dis- 
solved in  water,  its  molecular  weight  is  found  to  be  34, 
and  this  requires  us  to  write  the  formula  H202,  not  HO. 

"When  a  substance  can  neither  be  dissolved  nor  vaporized 
without  decomposition,  both  these  methods  of  determining 
the  molecular  weight  fail.  Such  a  substance  is  mercuric 
oxide.  At  present  the  question  whether  it  should  be  written 
HgO  or  Hg202,  or  some  higher  multiple,  is  one  to  which 
we  can  not  return  an  answer.  Therefore  we  use  the  simplest 
formula  for  the  present.  The  problem  of  determining  the 
molecular  weights  of  such  substances  is  now  being  attacked, 
and  the  result  so  far  is  that  no  multiplying  is  necessary; 
the  molecular  weight  obtained  corresponds  to  the  simplest 
formula  which  will  express  the  composition  accurately. 
There  is  no  reason,  therefore,  to  consider  the  molecules  of 
solids  as  large  and  heavy  compared  with  those  of  gases. 


CHAPTER   XVII 

ACIDS,  BASES,  AND  SALTS -ELECTROLYTIC  DISSOCIATION- 
METALS  AND  NON-METALS 

167.  Some  strong  solution  of  sodium  hydroxide  is  placed 
in  a  beaker  and  strong  hydrochloric  acid  is  carefully  added 
to  it.     There  is  an  energetic  reaction,  the  liquid  becomes 
very  hot,  and  a  white  powder  of  common  salt  separates : 

NaOH  +  HC1  =  NaCl  +  H20. 

A  few  drops  of  the  same  sodium  hydroxide  solution  are 
placed  in  a  test-tube,  and  sulphuric  acid  added,  one  drop  at 
a  time.  Each  drop  of  the  acid  hisses  when  it  strikes  the 
liquid  like  water  falling  upon  red-hot  iron.  Great  heat  is 
evolved,  and  white  crystals  of  sodium  sulphate  are  produced : 

2NaOH  +  H2S04  =  Na2S04  +  2H20. 

Sodium  hydroxide  is  a  base,  and  HC1  and  H2S04  are 
acids,  and  these  two  experiments  will  serve  as  illustrations 
of  what  happens  when  an  acid  and  a  base  come  into  contact. 
A  chemical  change  occurs.  Water  is  one  product  and  the 
other  is  a  substance  called  a  salt,  which  consists  of  the  acid, 
with  its  hydrogen  replaced  by  the  metal  of  the  base. 

168.  Bases. — The  bases  are  hydroxides  of  the  metals  or 
of  some  group  of  atoms — NH4,  for  instance — which  plays 
the  role  of  a  metal.    When  a  base  is  dissolved  in  water  the 
solution  has  a  bitter,  burning  taste  and  a  caustic  effect  upon 
the  skin  and  other  tissues,  which  is  stronger  the  stronger 
the  base  and  the  more  concentrated  the  solution.    There  are 
various  sensitive  coloring  matters  called  indicators,  whose 

125 


126  ELEMENTARY  CHEMISTRY 

colors  change  remarkably  when  a  base  acts  upon  them. 
Thus,  a  water  solution  of  red  litmus  is  turned  deep  blue  by 
a  little  sodium  hydroxide  or  other  base ;  a  solution  of  coch- 
ineal, orange  itself,  turns  violet. 

169.  Acids. — Acids  are  compounds  containing  hydrogen 
which  can  be  replaced  by  metals,  producing  salts.     Acids 
have  a  sour  taste — the  acid  taste — similar  to  that  of  vinegar 
or  lemon  juice,  and  strong  acids  have  an  energetic  caustic 
action    upon    organic    tissues.      Acids    reverse    the    color 
changes   which   bases   produce    with   indicators.      Litmus, 
which  has  been  turned  from  red  to  blue  by  a  base,  is  again 
colored  red  by  an  acid;  cochineal,  which  has  been  colored 
violet  by  a  base,  again  becomes  orange  when  an  acid  is 
added. 

170.  Neutralization. — The  reaction  between  an  acid  and 
a  base  in  which  a  salt  and  water  are  produced  is  called 
neutralization.     From  the  equation — 

NaOH  +  HC1  =  NaCl  +  H20, 
40  36.5        58.5          18, 

in  which  the  numbers  written  below  are  the  molecular 
weights  of  the  substances,  it  is  clear  that  if  hydrochloric 
acid  and  caustic  soda  are  mixed  in  exactly  the  right  quan- 
tities, the  liquid  will  contain  neither  substance;  it  will  con- 
tain  only  salt  and  water.  Thus,  if  we  dissolve  40  grams 
of  pure  sodium  hydroxide  in  water,  coloring  the  liquid  with 
a  little  litmus,  and  then  add  gradually  36.5  grams  of  hy- 
drochloric acid  dissolved  in  water,  the  litmus  will  remain 
blue  until  all  the  hydrochloric  acid  is  added.  At  this  point 
the  liquid  contains  nothing  but  sodium  chloride  and  water. 
One  drop  of  acid  in  excess  of  this  quantity  will  instantly 
change  the  color  to  red. 

171.  Salts. — A  salt  is  the  product  of  the  chemical  change 
between  an  acid  and  a  base.    It  is  a  compound  derived  from 
an  acid  by  replacing  the  hydrogen  by  a  metal,  e.  g.\ 


ANTOINE  LAURENT  LAVOISIER 
B.  Paris,  1743.     D.  on  the  scaffold,  1794. 


SALTS  127 

Sulphuric  acid  Sodium  sulphate  Calcium  sulphate 

H2S04.  Na2S04.  CaS04. 

Nitric  acid  Sodium  nitrate  Calcium  nitrate 

HN03.  NaN08.  Ca(N03)2. 

Salts  have  a  saline  taste  which  is  quite  different  from 
the  sour  taste  of  the  acids  and  the  bitter  taste  of  the  bases. 
A  solution  containing  only  a  salt  and  water  has  no  effect 
upon  the  color  of  indicators.  Some  salt  solutions  appear 
to  contradict  this  statement.  Thus,  if  we  dissolve  pure 
copper  chloride  in  water  the  liquid  reddens  litmus.  In  this 
case  the  solution  contains  hydrochloric  acid,  which  has  been 
produced  by  the  reaction  of  a  small  portion  of  the  copper 
chloride  with  water : 

CuCl2  +  2H20  =  Cu(OH)2  +  2HC1. 

On  the  other  hand,  when  potassium  cyanide,  KCN",  is 
dissolved  in  water,  the  solution  turns  red  litmus  blue.  This 
is  because  some  of  the  potassium  cyanide  has  reacted  with 
the  water: 

KCN  +  H20  =  KOH  +  HCN, 

and  the  liquid  contains  potassium  hydroxide,  an  extremely 
strong  base.  This  partial  decomposition  of  salts  by  water 
is  called  hydrolysis.  It  is  frequent  with  salts  of  weak  acids 
and  bases. 

172.  Acids  only  show  their  peculiar  properties  in  pres- 
ence of  water. — It  is  reasonable  to  ascribe  the  peculiar  qual- 
ities of  acids  to  the  hydrogen  which  they  all  contain.  But 
the  remarkable  properties  of  acids  are  only  manifested  in 
water  solution  (compare  p.  82).  Acids  free  from  water 
do  not  affect  indicators,  and  are  inert  toward  many  sub- 
stances upon  which  we  should  expect  them  to  act.  Thus, 
when  they  are  brought  into  contact  with  metals,  either  there 
is  no  reaction  or  the  reaction  is  quite  different  from  that 
which  takes  place  when  the  same  metal  is  treated  with  a 
water  solution  of  the  same  acid.  Further,  there  are  thou- 
sands of  substances  containing  hydrogen  which  are  not 


128  ELEMENTARY  CHEMISTRY 

acids.  Alcohol,  C2H60,  for  example,  when  dissolved  in 
water,  yields  a  solution  which  has  no  acid  taste,  does  not 
affect  the  color  of  indicators,  and  does  not  exhibit  any  of 
the  chemical  behavior  of  an  acid. 

Acids,  then,  differ  from  other  substances  in  this  respect, 
that  they  contain  hydrogen,  which,  in  water  solution,  is  in 
a  different  state  from  the  hydrogen  of  other  compounds. 
What  is  the  difference  ? 

173.  Solutions  of  acids  in  water  conduct  the  electric 
current. — A  solution  of  alcohol  in  water  is  an  insulator; 
it  does  not  allow  the  electric  current  to  pass.    Liquid,  water- 
free  hydrochloric  acid  behaves  in  the  same  way.     On  the 
other  hand,  the  current  is  freely  transmitted  by  a  water 
solution  of  hydrochloric  acid,  and  its  transmission  is  car- 
ried on  by  a  stream  of  hydrogen  particles  which  pass  to  the 
negative  pole,  give  up  their  electric  charges  to  the  electrode, 
and  are  liberated  as  gas,  and  a  procession  of  chlorine  par- 
ticles, which  move  to  the  positive  pole  and  are  there  re- 
lieved of  their  charges  and  liberated  in  the  same  way.     It 
appears,  then,  that  the  hydrogen  and  chlorine  atoms  in  a 
water  solution  of  hydrochloric  acid  exhibit  an  independence 
of  each  other,  a  readiness  to  move  in  different  directions, 
which  we  do  not  find  in  the  anhydrous  substance.    In  con- 
nection with  this  we  must  think  of  the  inertness  of  the  an- 
hydrous substance  and  the  great  chemical  activity  of  the 
water  solution.    This  activity  is  associated  with  the  ability 
to  conduct  the  current,  for  HC1  dissolved  in  chloroform  is 
a  non-conductor,  and  is  just  as  inactive  chemically  as  the 
anhydrous  substance. 

174.  Ions. — Facts  of  this  kind  have  led  chemists  to  pic- 
ture to  themselves  the  state  of  things  in  a  water  solution 
of  hydrochloric  acid  somewhat  in  this  way.    A  large  propor- 
tion of  the  acid  exists  no  longer  as  molecules  HC1,  but  as 
atoms  H  and  Cl.    Each  H  atom  has  a  positive  charge  and 
each  Cl  atom  a  negative  charge,  and  they  move  about  quite 


ELECTROLYTIC  DISSOCIATION  129 

independently  of  each  other.  When  the  electrodes  from  the 
battery  are  dipped  into  the  liquid  the  H  atoms  with  their 
positive  charges  naturally  move  to  the  negative  pole,  for 
the  same  reason  that  a  positively  charged  pith-ball  is  at- 
tracted by  a  negatively  charged  stick  of  sealing-wax.  Here 
they  yield  up  their  charges  and  the  discharged  atoms  unite 
in  pairs  to  molecules  of  ordinary  hydrogen  H2,  which  is 
liberated.  The  same  thing  happens  with  the  chlorine  at  the 
positive  pole.  These  charged  atoms  we  shall  call  ions,  and 
it  is  impossible  to  obtain  a  liquid  containing  hydrogen  ions 
or  chlorine  ions  alone,  because  no  positive  charge  can  exist 
anywhere  without  an  equivalent  negative  charge  in  the 
neighborhood,  and  vice  versa.  Although  the  hydrogen  and 
chlorine  in  a  water  solution  of  hydrochloric  acid  exist  largely 

as  separate  ions  H  and  Cl,  yet  we  must  not  expect  the  solu- 
tion to  exhibit  the  odor  and  bleaching  action  of  chlorine,  nor 
the  combustibility  and  other  properties  of  hydrogen.  For 
the  properties  of  ordinary  chlorine  are  those  of  a  substance 
composed  of  molecules  C12,  and  the  water  solution  of  hydro- 
chloric acid  does  not  contain  these  molecules.  It  contains 

the  ions  Cl  with  their  negative  charges — quite  a  different 
thing.  The  same  statement  applies  to  the  hydrogen. 

175.  Difference  between  strong  and  weak  acids. — The 
properties  of  acids,  as  a  class,  are  those  of  the  hydrogen 
ions,  and  the  strongest  acids  are  those  in  which  hydrogen 
ions  are  most  abundant — that  is,  those  in  which  the  separa- 
tion sketched  above  is  most  complete.  In  a  solution  of  a 
very  strong  acid,  like  HC1,  a  large  fraction  of  the  substance 
is  separated— almost  all  of  it  in  a  somewhat  dilute  solution 
— and  the  liquid  strongly  exhibits  those  peculiar  properties 
which  are  due  t'o  the  presence  of  hydrogen  ions — acid  prop- 
erties, we  call  them.  On  the  other  hand,  in  a  solution  of  a 
weak  acid  (like  HCN,  prussic  acid),  the  amount  of  sepa- 
ration is  slight.  Most  of  the  acid  still  exists  as  molecules 


130  ELEMENTARY  CHEMISTRY 

HCN,  the  number  of  hydrogen  ions  is  small,  and  the  acid 
properties  feeble.  Since  the  passage  of  the  electric  current 
through  the  liquid  is  carried  on  by  means  of  the  ions,  it 
follows  that  a  solution  of  a  strong  acid — one  very  much 
separated  into  ions — will  conduct  well,  while  a  solution  of 
a  weak  acid — one  in  which  the  separation  is  slight — will 
conduct  badly.  Hence,  the  strength  of  an  acid  can  be  meas- 
ured ~by  its  ability  to  conduct  the  current.  The  more 
marJced  the  acid  properties  the  more  readily  the  electric  cur- 
rent will  pass  through  the  water  solution  of  the  substance. 

176.  Solutions  of  bases  contain  hydroxyl  ions,  OH. — The 
same  views  apply  to  bases.    Since  the  bases  are  all  hydrox- 
ides, their  peculiar  properties  are  those  of  hydroxyl,  OH, 
and  in  solutions  of  the  bases  the  hydroxyl  exists  as  an  ion, 

OH.1    In  water  solutions  of  a  strong  base,  like  NaOH,  the 

+ 
separation   into   ions   (Na   and    OH)    is   almost   complete, 

the  hydroxyl  ions  are  numerous,  and  the  basic  properties 

strongly  in  evidence.    In  solutions  of  weak  bases  (NH4OH 

+ 
will  serve  as  an  instance),  the  separation  into  ions  (NH4 

and  OH)  is  slight,  the  hydroxyl  ions  few,  and  the  basic 
properties  feeble.  Finally,  there  are  many  hydroxides  (alco- 
hol, which  we  may  write  C2H5OH  this  time,  is  again  a  good 
example)  where  there  is  no  separation  into  ions  when  the 
substance  is  dissolved,  and  therefore  no  basic  properties. 
In  striking  agreement  with  this  is  the  fact  that  solutions 
of  strong  bases  are  good  conductors  of  the  electric  current, 
solutions  of  weak  bases  bad  conductors,  and  solutions  of  hy- 
droxides, like  alcohol,  which  are  not  bases  at  all,  non-con- 
ductors. 

177.  Ions  of  salts. — Salts  are  also  separated  into  ions 
when  dissolved  in  water,  and  therefore  their  solutions  con- 

1  We  shall  indicate  the  charges  of  the  ions  by  small   +  and  — 
signs  placed  above. 


ELECTROLYTIC  DISSOCIATION  131 

duct  the  current.    Nickel  chloride,  NiCl2,  for  instance,  exists 

in  solution  separated  into  two  chlorine  ions,  Cl  Cl,  and  one 
nickel  ion,  which  must  have  twice  as  much  positive  elec- 
tricity on  it  as  one  chlorine  ion  has  of  negative.  We  can 

+  + 
write  it  Ni.     All  solutions  of  nickel  salts  have  the  same 

green  color,  and  our  theory  of  ions  easily  explains  this  fact, 
for  this  color  is  simply  that  of  the  nickel  ion.  Otherwise  it 
would  be  astonishing  that  salts  so  different  in  composition 
as  nickel  chloride  (NiCl2),  nickel  sulphate  (NiS04),  and 
nickel  nitrate  (Ni(N03)2)  should  produce  exactly  the  same 
color.  When  the  current  is  passed  through  a  solution  of 
nickel  chloride,  the  chlorine  ions  proceed  to  the  positive  pole, 
where  they  yield  up  their  charges  and  are  liberated  as  chlo- 
rine gas,  just  as  in  the  case  of  hydrochloric  acid.  The  nickel 
ion  carries  its  positive  charge  to  the  negative  pole,  where  it 
gives  it  up  and  is  converted  into  ordinary  metallic  nickel, 
which  adheres  to  the  electrode.  Table  salt  in  solution  is 

+ 
separated  into  ions  Na  and  Cl,  and  when  the  current  passes 

through  it,  the  chlorine  is  liberated  at  the  positive  pole.  But 
it  is  impossible  for  sodium  to  be  liberated  at  the  other  pole 
on  account  of  the  presence  of  water.  Hence,  at  the  nega- 
tive pole,  hydrogen  from  the  water  escapes  instead. 

178.  Acids,  bases,  and  salts  are  the  only  substances 
whose  solutions  conduct  the  current,  and  the  only  sub- 
stances, therefore,  which  separate  into  ions  when  dissolved. 

An  ion  may  be  an  atom  or  a  group  of  atoms.     Thus,  the 

+ 
ions  of  hydrochloric  acid  are  simply  H  and  Cl,  but  those 

of  nitric  acid  are  H  and  N03,  and  those  of  sulphuric  acid 

(H2S04)  are  H,  H,  and  S04. 

This  idea  of  the  separation  of  dissolved  substances  into 
ions  has  thrown  a  flood  of  light  upon  chemical  and  physical 
changes  in  which  acids,  bases,  or  salts  are  concerned.    It  is 
10 


132  ELEMENTARY  CHEMISTRY 

called  the  theory  of  electrolytic  dissociation.  The  student 
who  pursues  chemical  studies  further  will  meet  with  it 
again. 

179.  Metals  and  non-metals. — When  we  compare  ele- 
ments like  copper  and  silver,  on  the  one  hand,  with  elements 
like  sulphur  and  phosphorus,  on  the  other,  certain  differ- 
ences force  themselves  on  our  attention.  The  most  notice- 
able distinction  is  the  appearance.  A  polished  surface  of 
copper  or  silver  reflects  almost  all  the  light  which  falls  upon 
it,  and  this  gives  to  it  what  is  called  the  metallic  luster. 
Again,  a  piece  of  silver  when  struck  heavily  with  a  hammer 
flattens  out,  and  with  care  can  be  beaten  into  a  thin  sheet ; 
a  piece  of  sulphur  flies  to  pieces.  The  silver  is  malleable, 
the  sulphur  brittle. 

Elements  like  copper  and  silver  are  called  metals;  ele- 
ments like  sulphur  and  phosphorus  non-metals. 

0.  Physical  properties  of  metals. — The  metals  have  a 
peculiar  luster — the  metallic  luster.  They  are  far  better 
conductors  of  heat  and  of  the  electric  current  than  the  non- 
metals.  The  metals  can  be  drawn  into  wire  and  beaten  out 
with  the  hammer — that  is,  they  are  ductile  and  malleable. 
Gold  can  be  beaten  into  leaf  so  thin  that  a  pile  containing 
a  hundred  thousand  sheets  in  contact  would  be  less  than 
one  centimeter  in  height,  and  one  gram  of  gold  can  be 
drawn  into  a  wire  three  thousand  meters  in  length.  All  of 
the  other  metals  are  inferior  to  gold  in  these  properties, 
and  in  some — lead  and  bismuth,  for  instance — malleability 
and  ductility  are  slight.  These  two  properties,  together 
with  the  tenacity  (toughness)  of  the  metals,  are  the  chief 
source  of  their  great  importance  in  commerce.  The  non- 
metals  are  immensely  behind  the  metals  in  all  three  re- 
spects. 

181.  Chemical  distinctions. — The  hydrogen  compounds 
of  the  non-metals  are  gases — water,  which  is  a  liquid  easily 
converted  into  a  vapor,  is  the  chief  exception — and  they  are 


METALS  AND   NON-METALS  133 

rather  stable  substances  which  show  no  very  marked  tend- 
ency to  separate  into  their  constituents.  Ammonia  and 
hydrochloric  acid  are  good  examples.  On  the  other  hand, 
many  of  the  metals  do  not  form  hydrogen  compounds  at 
all,  and  when  they  do,  the  substances  are  solids  and  decom- 
pose readily  into  hydrogen  and  the  metal. 

The  hydroxide  of  a  metal  is  a  base,  and  when  it 
dissolves  in  water  it  separates  into  positively  charged 
metal  ions  and  negatively  charged  hydroxyl  ions.  Thus, 

a   water   solution   of  potassium    hydroxide    (KOH)    con- 

+ 
tains  potassium  ions  positively  charged  (K)  and  hydroxyl 

ions  charged  negatively  (OH).  A  solution  of  calcium  hy- 
droxide, Ca(OH)2,  consists  of  calcium  ions  with  a  double 

positive  charge  (Ca)  and  twice  as  many  hydroxyl  ions,  each 

with  its  negative  charge  {_     J. 

\OH/ 

Since  the  solutions  of  the  metallic  hydroxides  contain 
hydroxyl  ions,  they  turn  red  litmus  blue  and  exhibit  the 
other  properties  which  we  have  learned  to  associate  with 
bases  (p.  125).  Many  of  the  metallic  hydroxides  are  insol- 
uble in  water — nickel  and  copper  hydroxides,  for  example — 
and  these  can  not  act  upon  the  indicators ;  but  they  can  still 
be  considered  as  bases  in  the  sense  that  they  react  readily 
with  acids,  forming  salts.' 

182.  Metals  tend  to  form  positively  charged  ions. — This 
tendency  of  the  metals  to  exist  in  solution  as  positively 
charged  ions  is  by  no  means  restricted  to  the  hydroxides. 

We  have  seen  that  in  a  solution  of  table  salt  there  are 

+ 
sodium  ions  (Na)  and  chlorine  ions  (Cl),  and  in  the  water 

solution  of  every  salt  the  metal  is  present  in  this  condition. 
A  metal  never  exists  alone  as  a  negatively  charged  ion  in 
solution,  and  therefore  never  proceeds  alone  to  the  positive 
pole  when  the  current  is  passed  through  the  liquid — always 


134  ELEMENTARY   CHEMISTRY 

to  the  negative  pole.  We  may  say,  therefore,  that  a  metal 
is  a  substance  which  has  a  tendency  to  exist  alone  in  solu- 
tion as  a  positively  charged  ion.  It  must  not  be  inferred 
from  the  word  "  alone  "  in  the  above  statement  that  there 
can  be  such  a  thing  as  a  solution  containing  positively 
charged  metal  ions  without  any  corresponding  negative  ions. 
This  is  impossible,  for  in  every  solution  the  number  of 
positive  and  negative  charges  must  be  exactly  the  same. 
We  mean  simply  that  the  metal  atom  moves  about  inde- 
pendently in  the  liquid  with  its  positive  charge,  and  pro- 
ceeds alone  to  the  negative  pole  in  electrolysis.  This  be- 
havior we  can  sum  up  by  saying  that  the  metals  are  electro- 
positive. 

183.  The  hydroxides  of  the  non-metals  are  acids. — Many 
of  the  hydroxides  of  the  non-metals  are  unstable  or  un- 
known, but  those  which  exist  are  acids ;  and,  since  the  proper- 
ties of  an  acid  liquid  are  simply  those  of  the  hydrogen  ions, 
they  must  separate  when  dissolved  in  quite  a  different  way 
from  the  metallic  hydroxides.  Chlorine  hydroxide  (C10H), 

which  is  called  hypochlorous  acid,  is  an  example.    Its  sepa- 

+ 
ration  when  dissolved  is  into  hydrogen  ions  (H)  and  ions 

(CIO).  An  atom  of  a  non-metal  can  never  exist  alone  as  a 
positively  charged  ion,  only  as  a  negatively  charged  one,  and 
therefore  can  never  proceed  by  itself  to  the  negative  pole 
when  the  current  is  passed — only  to  the  positive.  We  can 
sum  up  this  behavior  in  the  statement  that  the  non-metals 
are  electronegative.  It  follows  from  this  that  there  can  be 
no  such  thing  as  a  salt  in  which  a  single  non-metallic  atom 
is  the  positive  constituent,  no  such  thing  as  a  nitrate  of 
chlorine,  C1N03,  or  a  sulphate  of  bromine,  Br2S04.  Such 
a  compound  is  a  chemical  impossibility.  Only  the  metals 
form  salts  with  acids. 


CHAPTEE    XVIII 

THE    SODIUM    GROUP 

This  group  contains  the  following  five  elements : 

Lithium,  Li.  ^   Rubidium,  Rb.    3  ^          /.£"  i- 

Sodium,  Na.    q  >.          \  f^  Caesium,  Cs.         i.  i,*?'  />  §£* 
Potassium,  K.      :  .5**  %^ 

184.  General  properties. — This  group  contains,  from  the 
chemical    standpoint,    the    most    positive    metals    known. 
Much  of  what  we  have  learned  about  sodium  is  true  of 
all  five.    They  are  soft  enough  to  be  cut  with  a  knife,  are 
excellent  conductors  of  heat  and  electricity,  and  possess 
small  densities.     Lithium,  the  lightest,  is  little  more  than 
half  as  dense  as  water,  and,  except  solid  hydrogen,  is  the 
lightest  of  all  solids.     These  elements  possess  a  strong  sil- 
ver-white metallic  luster  which  is  permanent  when  the  spec- 
imen is  sealed  up  in  dry  hydrogen,  but  rapidly  disappears 
in  the  air.     This  is  due  to  the  oxygen  of  the  air,  which 
acts  upon  the  surface,  producing  the  oxide  of  the  metal,  and 
still  more  to  the  water-vapor,  which  produces  the  hydroxide. 

185.  Reaction  with  water. — When  one  of  these  metals 
comes  in  contact  with  water,  hydrogen  is  given  off  and  the 
hydroxide  of  the  metal  dissolves  in  the  excess  of  water. 
We  have  studied  this  action  in  the  case  of  sodium: 

Na  +  H20  =  NaOH  +  H. 

The  same  change  occurs  with  the  others.  If  we  indicate 
an  atom  of  any  one  of  the  five  by  M,  we  can  write  an 
equation — 

M  +  H20  =  MOH  +  H, 

135 


136  ELEMENTARY  CHEMISTRY 

which  describes  their  conduct  with  water,  and  the  equation 
for  a  particular  metal  is  obtained  by  substituting  the  sym- 
bol of  that  metal  for  M  on  both  sides  of  the  expression. 

The  energy  of  the  reaction  with  water  increases  with 
increasing  atomic  weight.  Lithium  evolves  hydrogen  rap- 
idly, but  does  not  melt.  Sodium  melts  to  a  globule  which 
runs  about  on  the  surface  of  the  water,  but  the  temperature 
does  not  rise  high  enough  to  inflame  the  hydrogen.  In  the 
case  of  potassium  the  hydrogen  catches  fire.  Rubidium 
and  caesium  react  with  water  with  explosive  violence. 

186.  Occurrence. — Of  course  elements  like  those  of  this 
group — energetically  acted  upon  by  oxygen  and  water — 
are   not   found   free   in   nature.      Compounds   of   sodium 
and  potassium   are   abundant,   the   other   three   are   rare. 
Lithium,  though  it  nowhere  occurs  in  large  quantities,  is 
quite  widely  distributed.    Traces  of  it  are  contained  in  many 
soils,  in  sea-water  and  most  natural  waters,  in  plants  (for 
instance,  in  tobacco,  coffee,  sugar-cane,  and  seaweed),  and 
even  in  substances  of  animal  origin  (in  blood  and  milk,  for 
example).    This  wide  distribution  is  the  rule  with  elements 
of  small  atomic  weight  (Li  =  7). 

187.  Tendency  to  form  positive  ions. — In  the  violence 
with  which  the  metals  of  this  group  react  with  water — 
forming  their  hydroxides — we  see  the  strong  tendency  of 
their  atoms  to  assume  the  state  of  positively  charged  ions 
in  solution,  and  this  tendency  is  greater  in  this  group  than 
anywhere  else.     This  is  another  and  better  way  of  saying 
that  in  these  elements  the  metallic  properties  are  most 
highly  developed.    We  shall  not  be  surprised,  therefore,  by 
the  fact  that  the  hydroxides  of  these  metals  are  the  strongest 
bases  known  to  chemical  science,  for  this  tendency  of  the 
metal  atom  will  cause  the  hydroxide  to  cleave  very  com- 
pletely into  positively  charged  metal  ions  and  negatively 
charged  hydroxyl  ions  when  dissolved  in  water.     Hydroxyl 
ions  will  therefore  be  numerous  in  the  liquid,  and  the  char- 


THE  SODIUM  GROUP  137 

acteristic  properties  which  they  communicate  to  a  solution 
will  be  strongly  marked — in  other  words,  the  dissolved  sub- 
stance will  be  a  strong  *base.  In  this  case,  as  in  most  other 
groups  composed  of  metals,  the  strength  of  the  base  in- 
creases with  the  atomic  weight  of  the  metal.  Cesium  hy- 
droxide, CsOH,  is  the  strongest  of  all  bases. 

188.  Salts  containing  the  metals  of  this  group  are  al- 
most all  freely  soluble  in  water.     The  ions  of  all  five  ele- 
ments are  colorless,  and  accordingly  solutions  containing 
salts  of  these  metals  are  colorless  also,  unless  the  acid  ion 
possesses  a  color. 

189.  Lithium,  Li  =  7.— Lithium  is  a  silver- white  metal,  harder 
than  sodium.     It  oxidizes  readily  in  the  air,  but  not  so  rapidly  as 
the  other  members  of  the  subgroup.     Thrown  into  water,  it  is  rap- 
idly converted  into  its  hydroxide — 

Li  +  H9O  =  LiOH  +  H— 

but  the  temperature  does  not  rise  high  enough  to  melt  the  metal. 
Heated  in  the  air  or  in  oxygen,  it  burns  with  intense  white  light  to 
lithium  oxide,  Li2O. 

Lithium  chloride,  LiCl,  is  white,  deliquescent,  and  very  soluble 
in  water.  Some  mineral  waters  contain  considerable  quantities  of 
it,  and  such  waters  are  employed  with  advantage  in  rheumatic 
troubles. 

190.  Valence. — An  atom  of  lithium  is  equal  in  combin- 
ing power  to  an  atom  of  hydrogen.    Thus  it  replaces  hydro- 
gen atom  for  atom  in  water : 

Water  Lithium  hydroxide 

HOH,  LiOH, 

in  hydrochloric  acid: 

Hydrochloric  acid  Lithium  chloride 

HC1,  LiCl, 

and  in  all  other  acids.  The  combining  power  of  an  ele- 
ment has  received  the  name  of  valence,  and  if  we  call  the 
combining  power  of  hydrogen  one,  we  can  describe  this  state 
of  things  by  the  statement  that  lithium  also  has  a  valence 


138  ELEMENTARY  CHEMISTRY 

of  one,  or  is  univalent.  The  great  advantage  of  putting 
the  matter  in  this  way  is  that  it  enables  us  to  write  at 
once  the  formula  of  any  lithium  salt  if  we  know  that  of 
the  corresponding  acid.  All  that  is  necessary  is  to  replace 
each  hydrogen  atom  by  an  atom  of  lithium.  Thus : 

Sulphuric  acid  Lithium  sulphate 

H2S04.  Li2S04. 

Nitric  acid  Lithium  nitrate 

HN03.  LiN03. 

The  five  metals  of  this  group  are  univalent  in  almost 
all  their  compounds;  and  if  we  remember  this,  it  is  un- 
necessary to  memorize  the  formulas  of  their  salts,  for  we  can 
write  them  in  the  same  way  as  we  have  done  with  lithium 
sulphate  and  nitrate.  Thus,  the  idea  of  valence  is  a  valuable 
labor-saving  device  in  dealing  with  elements  like  these, 
whose  valence  remains  the  same.  The  atoms  of  some  ele- 
ments, like  nitrogen,  vary  in  combining  power  in  different 
compounds,  and  then  the  advantage  of  the  notion  of  valence 
is  not  so  great.  We  are  compelled  to  learn  the  formula  of 
each  compound  separately. 

Sodium,  the  second  member  of  the  group,  has  already 
been  described  (Chapter  IX). 

POTASSIUM,  K  =  39. 

191.  Occurrence. — Potassium  is  the  most  abundant  ele- 
ment of  the  group.  A  reference  to  the  table  in  the  Ap- 
pendix will  show  that  the  accessible  parts  of  the  earth's 
crust  contain  on  an  average  2.4  per  cent  of  this  element. 
Almost  all  the  common  rocks — granite,  gneiss,  basalt,  and 
so  on — contain  potassium  compounds ;  and  when  they  break 
up  under  the  influence  of  the  weather,  the  potassium  com- 
pounds pass  into  the  soil,  which  always  contains  them. 
From  the  soil  they  are  absorbed  by  plants:  potassium  is 
indispensable  to  vegetable  life. 


THE  SODIUM  GROUP  139 

192.  Importance  of  potassium  in  the  life  process. — Many 
experiments  have  shown  that  the  place  of  the  potassium 
in  plants  can  not  be  taken  by  the   other   metals   of  the 
group,  for  when  the  attempt  is  made  to  displace  the  potas- 
sium by  watering  the  plant  systematically  with  a  solution 
of  a  lithium  compound,  for  instance,  the  plant  shows  signs 
of  poisoning,  and  finally  dies.     From  the  plants  potassium 
finds  its  way  into  the  bodies  of  animals,  and  it  is  contained 
in  all  animal  tissues,  even  in  muscular  tissue,  which  does 
not  contain  sodium. 

193.  History  and  preparation. — When  a  plant  is  burned, 
the  potassium  is  left  in  the  ashes  as  potassium  carbonate, 
K2C03,  and  this  can  be  separated  from  the  other  substances 
present  by  treating  the  ashes  with  water,  in  which  potas- 
sium carbonate  dissolves.     Wood-ashes  were  the  earliest 
source  of  potassium  compounds,  but  they  have  ceased  to  be 
important  on  account  of  the  discovery  of  immense  deposits 
of  potassium  chloride,  KC1,  at  Stassfurt,  in  Germany. 

Potassium  was  first  obtained  almost  at  the  same  time 
as  sodium  by  Sir  Humphry  Davy  (p.  68),  who  decomposed 
potassium  hydroxide,  KOH,  by  the  electric  current.  It  is 
made  by  heating  potassium  carbonate  to  whiteness  with 
powdered  charcoal: 

K2C03  +  20  =  3CO  +  2K. 

Potassium  is  liberated  as  a  green  vapor,  which  is  con- 
densed under  petroleum.  The  process  is  dangerous  and 
explosions  are  frequent.  Potassium  has  ceased  to  be  an 
important  product,  because  sodium,  which  can  be  easily  and 
cheaply  made  by  the  electric  process  (p.  68),  answers  the 
same  purpose  in  almost  all  technical  operations. 

194.  Properties. — Potassium  is  a  metal,  with  a  silver- 
white   metallic  luster  which  disappears   in  the   air   more 
rapidly  than  that  of  sodium.    Heated  in  the  air,  or  in  gases 
containing  oxygen,  it  melts  to  a  mercury-like  liquid,  which 


140  ELEMENTARY  CHEMISTRY 

takes  fire  and  burns  with  a  violet  flame  to  gray  potassium 
oxide,  K20,  or  to  yellow  potassium  tetr oxide,  K20±,  or  to  a 
mixture  of  both,  according  to  the  temperature  and  the 
amount  of  oxygen  supplied  to  it.  It  decomposes  water  en- 
ergetically, thus: 

K  +  H20  =  KOH  +  H, 

and  the  hydrogen  evolved  catches  fire  at  once,  burning  with 
a  flame  colored  violet  by  potassium. 

195.  Potassium  hydroxide,  KOH,  commonly  called  caus- 
tic potash,  was  formerly  made  in  the  same  way  as  sodium 
hydroxide   (p.   70),  by  boiling  potassium  carbonate  with 
slaked  lime: 

Ca(OH)2  +  K2C03  =  CaC03  +  2KOH. 

At  present  it  is  prepared  by  passing  the  electric  current 
through  a  solution  of  potassium  chloride.  This  solution 

contains  K  ions  positively  charged  and  Cl  ions  charged  nega- 
tively. When  the  current  passes  the  chlorine  ions  proceed 
to  the  positive  pole,  give  up  their  charges,  and  are  liberated 
as  chlorine  gas,  which  is  used  for  the  manufacture  of  bleach- 
ing-powder  (p.  88).  The  potassium  ions  go  to  the  negative 
pole,  but  of  course  a  metal  like  potassium  can  not  be  set  free 
in  the  presence  of  water.  Instead,  hydrogen  is  liberated 
from  the  water.  This  leaves  hydroxyl  ions;  and  since  the 
liquid  around  the  negative  pole  contains  potassium  ions  also, 
it  is  simply  a  solution  of  potassium  hydroxide,  which  can  be 
obtained  solid  by  evaporation.  Chlorine  and  potassium 
hydroxide  act  upon  each  other  chemically  (p.  87),  and  there- 
fore it  is  necessary  to  keep  them  apart.  This  is  accom- 
plished by  placing  between  the  positive  and  negative  poles  a 
porous  partition,  which  allows  the  current  to  pass,  but  does 
not  permit  the  liquids  on  the  two  sides  to  mix. 

196.  Properties  of  potassium  hydroxide. — Potassium  hy- 
droxide is  very  similar  to  sodium  hydroxide.    It  usually  oc- 


THE  SODIUM  GROUP  141 

curs  in  commerce  in  the  form  of  sticks.  It  is  a  white  solid, 
melting  at  a  red  heat  and  vaporizing  at  higher  temperatures 
without  decomposition.  It  is  excessively  soluble  in  water, 
and  the  solution  acts  energetically  upon  vegetable  and  ani- 
mal tissues,  and  possesses  in  a  marked  degree  all  those 
properties  which  we  have  learned  to  associate  with  the  pres- 
ence of  hydroxyl  ions. 

197.  Compounds  of  potassium  with  sulphur. — When  po- 
tassium is  gently  heated  with  sulphur,  the  two  combine  with 
explosive  energy,  and  the  result  is  a  fused  brownish-yellow 
mass,  which  may,  according  to  the  proportions  in  which 
the  elements  were  taken,  contain  any  or  all  of  the  following 
five  compounds: 

Potassium  monosulphide,  K2S; 
"         disulphide,  K2S2; 
"         trisulphide,  K2S3 ; 

tetrasulphide,  K2S4; 
"         pentasulphide,  K2S5. 

It  is  not  possible  to  obtain  any  of  these  compounds  in 
the  pure  state  by  heating  the  two  elements  together,  but 
they  have  all  been  prepared  by  other  methods.  They  are 
yellow  or  red  solids,  soluble  in  water. 

Sulphur  in  the  sulphides  has  a  combining  power  equal  to 
that  of  two  hydrogen  atoms.  This  is  shown,  for  example,  in 
hydrogen  sulphide,  H2S.  This  we  express  by  the  statement 
that  sulphur  in  these  compounds  has  a  valence- of  two,  or  is 
bivalent.  We  have  seen  that  the. metals  of  the  sodium  group 
are  usually  univalent.  The  composition  of  potassium 
monosulphide  is  just  what  we  should  expect.  We  may  think 
of  it  as  hydrogen  sulphide,  H2S,  in  which  each  hydrogen 
atom  has  been  replaced  by  an  atom  of  potassium — 

H2S,  hydrogen  sulphide; 
K2S,  potassium  sulphide; 


142  ELEMENTARY  CHEMISTRY 

and  in  it  the  potassium  is  univalent,  the  sulphur  bivalent. 
But  what  are  we  to  think  of  the  other  compounds — of  K2S2, 
for  instance  ?  If  each  sulphur  atom  is  bivalent  in  this  sub- 
stance, the  two  must  have  a  combining  power  of  four;  and 
since  it  only  requires  two  potassium  atoms  to  combine  with 
them,  the  potassium  must  be  bivalent.  Similar  reasoning 
shows  us  that  if  the  sulphur  is  always  bivalent,  then — 

In  KjSs  the  potassium  has  a  valence  of  three  or  is  trivalent. 
InK8S4  "          "  "  "        four      "     quadrivalent. 

InK,S6  "          "  "  "        five       "     quinquivalent. 

So  that  we  seem  forced  to  admit  that,  while  potassium  is 
almost  always  univalent,  yet  in  its  compounds  with  sulphur 
and  in  some  other  compounds — K204,  for  instance — it  may 
vary  in  combining  power  from  one  to  five. 

This  conclusion,  that  the  same  atom  may  have  different 
combining  powers  at  different  times,  is  unpleasant  to  some 
chemists.  Here  is  a  way  of  escape  from  it.  Let  us  admit 
that  in  K2S2  the  sulphur  atoms  are  united  together,  and  each 
holds  a  potassium  atom,  thus : 

K—S— S— K, 

then  we  can  write  a  formula  for  the  compound  and  still 
believe  that  the  potassium  is  univalent.  In  the  same  way 
potassium  trisulphide,  K2S3,  can  be  written — 

K — S — S — S — K , 
potassium  tetrasulphide,  K2S4 — 

K— S— S— S— S— K; 
and  potassium  pentasulphide,  K2S5— 

K— S— S— S— S— S— K. 

Formulas  like  these,  which  express  our  belief  regarding 
the  way  in  which  the  atoms  are  united  in  compounds,  are 
called  structural  formulas.  The  structural  formulas  of 
these  compounds  of  potassium  and  sulphur  are  of  no  im- 


THE  SODIUM  GROUP  143 

portance,  because  there  is  no  experimental  evidence  in  their 
favor.  But  the  student  should  carefully  notice  the  impor- 
tant idea  that  it  is  possible  to  think  of  the  molecule  of  a 
compound  as  a  structure  in  which  the  atoms  have  a  per- 
fectly definite  arrangement.  This  "  doctrine  of  chemical 
structure/'  as  it  is  called,  has  been  of  priceless  value  in 
dealing  with  the  innumerable  compounds  which  carbon 
forms  with  other  elements. 

198.  Potassium    chloride,    KC1.  —  When    potassium    is 
heated  gently  in  chlorine,  it  burns  with  a  violet-red  flame 
to  potassium  chloride,  KC1.    This  compound  occurs  in  great 
quantities  at  Stassfurt,  in  Germany,  and  it  is  the  raw  mate- 
rial  for   the   manufacture    of   all   potassium   compounds. 
Potassium  chloride  strongly  resembles  sodium  chloride.    It 
is  colorless  when  pure   and  crystallizes   in   cubes.     It  is 
quite  soluble  in  water.     Its  taste  resembles  that  of  salt, 
but  is  sharper.     The  fact  that  it  takes  the  place  of  salt 
in  the  diet  of  certain  savages  has  been  referred  to  (p.  67). 

199.  Potassium   bromide,  KBr,  and  potassium  iodide, 
KI,  are  white,  crystallize  usually  in  cubes,  and  are  readily 
soluble  in  water.     They  are  extensively  used  in  medicine 
and,  especially  the  first,  in  photography. 


r?  ^,f 

;o.r 

•/f. 


;o.r 


CHAPTER    XIX 

THE  COPPER  GROUP 
Copper,  Cu.  Silver,  Ag.  Gold,  Au. 

200.  General  properties. — The  metals  of  this  group  are 
not  so  strongly  metallic,  chemically,  as  the  metals  of  the 
sodium  group.  Silver  and  gold  are  not  affected  by  water 
or  by  oxygen  at  any  temperature,  nor  is  copper  affected  by 
either  in  the  cold.  Oxygen,  however,  combines  with  heated 
copper,  and  when  steam  is  led  over  melted  copper  at  a  white 
heat,  copper  oxide  is  formed  and  hydrogen  liberated,  but 
very  slowly. 

This  indifference  toward  water  and  oxygen  explains  the 
fact  that  these  metals  occur  free  in  nature,  and  therefore 
became  known  to  man  very  early,  while  the  elements  of  the 
Bodium  group  have  all  been  discovered  within  the  last  hun- 
dred years.  The  elements  of  the  sodium  group  are  not 
employed  for  any  of  the  purposes  which  we  associate  with 
metals,  for  their  conduct  toward  air  and  water  forbids  it. 
But  the  elements  of  the  copper  group  are  particularly 
adapted  to  such  uses — for  coinage,  for  example. 

Although  copper,  silver,  and  gold  are  far  less  metallic 
chemically  than  the  sodium  group,  they  possess  the  physical 
properties  of  a  metal  to  a  much  greater  extent. 

Copper,  silver,  and  gold  greatly  surpass  sodium  and  its 
allies  in  tenacity,  malleability,  and  ductility.  In  this  group 
tenacity  decreases  with  increasing  atomic  weight,  copper 
being  the  most  tenacious.  With  malleability  and  ductility 
the  reverse  is  true,  gold  being  by  far  the  most  malleable  and 
144 


THE  COPPER  GROUP  145 

ductile.  All  three  are  excellent  conductors  of  heat  and  of 
the  electric  current.  They  are  far  denser  than  the  ele- 
ments of  the  sodium  group — copper,  the  lightest,  being 
about  nine  times  as  dense  as  water — and  melt  at  higher 
temperatures,  at  a  bright  yellow  heat  or  beyond. 

COPPER,  Cu  =  63.5. 

201.  Occurrence. — Copper  is  widely  distributed  in  na- 
ture.    The  most  important  copper-mining  districts  are  in 
Montana,  Arizona,  and  Michigan,  where  great  quantities  of 
it  are  mined.    Many  compounds  of  the  metal — for  instance, 
the    oxide    Cu20    and   the    sulphide    Cu2S — are    abundant 
enough  to  be  important  as  ores.    Traces  of  copper  occur  in 
many  plants — hops  always  contain  it — and  in  the  animal 
kingdom  also — for  example,  in  the  red  feathers  of  certain 
birds  and  in  the  human  kidneys. 

202.  Physical  properties. — Pure  copper  has  a  rose-red 
color,  different  from  that  of  ordinary  copper.     It  is  very 
ductile,  malleable,  and  tenacious,  and  its  ability  to  conduct 
heat  and  the  electric  current  is  greater  than  that  of  any 
other  metal  except  silver.    It  melts  at  a  yellow  heat,  and  at 
the  high  temperature  of  the  electric  arc  (3500°)  it  rapidly 
boils  away,  torrents  of  brown  smoke  being  produced  by 
the  combination  of  the  vapor  to  CuO,  with  the  oxygen  of 
the  air. 

203.  Chemical  properties. — When  a   compact  piece   of 
copper  is  heated  in  the  air  it  combines  slowly  with  oxygen. 
Two  oxides  are  produced.    Next  the  copper  is  red  cuprous 
oxide,  Cu20,  and  outside  a  black  coating  of  cupric  oxide, 
CuO.    At  a  white  heat  melted  copper  burns  in  the  air  with  a 
green  light  to  cupric  oxide,  and  slowly  takes  the  oxygen 
from  steam  led  over  it,  thus : 

H20  +  Cu  =  CuO  +  H2. 

Air  and  water  at  ordinary  temperatures  affect  it  very  little, 
but  in  moist  air  containing  carbon  dioxide  it  becomes  cov- 


146  ELEMENTARY  CHEMISTRY 

ered  with  a  green  film  which  contains  the  hydroxide  and  the 
carbonate,  Cu(OH)2CuC03.  This  coating  is  called  verdi- 
gris. Copper  is  hardly  affected  by  hydrochloric  acid;  sul- 
phuric acid,  when  strong  and  hot,  attacks  it  rapidly,  and 
nitric  acid,  even  dilute  and  cold,  quickly  dissolves  it. 

204.  Action  upon  the  body. — Copper  and  copper  com- 
pounds are  not  very  poisonous,  and  when  they  are  taken 
into  the  system  daily  through  long  periods,  the  metal  does 
not  accumulate  in  the  body  and  finally  produce  poisoning, 
as  lead  does,  but  is  excreted  as  fast  as  introduced. 

205.  Uses. — Copper  is  an  important  metal  practically. 
Large  quantities  of  it  are  made  into  wires  for  the  pur- 
pose   of   conducting   the    electric    current,    and    it    makes 
part   of  many  important   alloys,  like   brass,  which   is   an 
alloy  of  copper  and  zinc,  and  bronze,  an  alloy  of  copper 
and  tin. 

206.  Compounds  of  copper  and  oxygen. — Six  compounds 
of  these  two  elements  are  known,  of  which  only  two  are 
important  enough  to  be  discussed  here.    These  are  cuprous 
oxide,  Cu20,  and  cupric  oxide,  CuO. 

It  will  be  noticed  that  if  oxygen  is  bivalent,  then  in  cupric  oxide 
the  copper  is  bivalent.  The  compound  can  be  viewed  as  water  in 
which  the  two  hydrogen  atoms  are  replaced  by  an  atom  of  copper : 

H30        CuO. 
With  cuprous  oxide  there  are  two  possibilities.     We  can  write  it 

Cu     Q 
Cu>0» 

in  which  case  it  is  water  in  which  two  hydrogen  atoms  are  replaced 
by  two  atoms  of  copper,  and  the  copper  is  univalent.  Or  we  can 
consider  that  the  copper  atoms  are  combined  together  and  also  with 
oxygen— 

Cux 

I    >0; 

Cu/ 

and  in  this  case  the  copper  will  be  bivalent,  just  as  it  is  in  cupric 
oxide.  Now,  so  far  as  the  oxides  are  concerned,  there  is  no  evidence 
in  favor  of  either  formula.  But  corresponding  to  each  oxide  is  a 


THE  COPPEU  GROUP 

whole  series  of  compounds.    Thus  hydrochloric  acid  converts  cupric 
oxide  into  cupric  chloride,  CuCla — 

CuO  +  2HC1  =  CuCla  +  HaO— 
and  cuprous  oxide  into  cuprous  chloride,  CuaCla — 
CuaO  +  2HC1  =  CuaCla  +  H3O. 

The  composition  of  cuprous  chloride  would  be  equally  well  de- 
scribed by  the  formula  CuCl,  but  it  has  been  converted  into  vapor, 
and  the  density  of  this  vapor  is  99.1  referred  to  hydrogen.  This 
shows  (Chapter  XVI)  that  the  molecular  weight  must  be  99.1  x  2, 
or  198.2,  and  therefore  we  must  write  the  formula  CuaCla — not 
CuCl — and  the  atoms  must  be  arranged  thus : 

Cu— Cl 


cl 


M— Cl. 

It  is  convenient  to  think  that  all  of  the  cuprous  compounds  contain 
a  pair  of  bivalent  copper  atoms  linked  together,  the  remaining  two 
valences  being  used  to  hold  the  other  elements  of  the  compound, 

Cu— 

cU 

207.  Cuprous  oxide,  Cu20,  occurs  in  nature,  and  is  called 
red  copper  ore.     It  can  be  made  by  heating  a  mixture  of 
cupric  oxide  and  finely  divided  copper,  air  being  excluded : 

Cu  +  CuO  =  Cu20. 

It  is  a  carmine-red  powder  insoluble  in  water.  As  we 
have  seen  above,  hydrochloric  acid  converts  it  into  white 
cuprous  chloride.  But  when  it  is  treated  with  sulphuric 
acid,  the  color  changes  to  the  rose-red  of  metallic  copper  and 
blue  cupric  sulphate  dissolves,  thus : 

Cu20  +  H2S04  =  Cu  +  CuS04  +  H20. 
Most  acids  behave  like  sulphuric  acid,  converting  it  into 
a  cupric  salt,  which  dissolves,  and  copper,  which  remains. 

208.  Cupric  oxide,  CuO,  commonly  called  copper  oxide, 
can  be  obtained  by  heating  copper  for  a  long  time  to  red- 
ness in  the  air,  or  by  heating  cupric  nitrate — 

Cu(NOs)2  =  CuO  +  2N02  +  0, 

the  reaction  being  similar  to  that  which  occurs  when  lead 
11 


148  ELEMENTARY   CHEMISTRY 

nitrate  is  heated  (p.  113).  It  is  a  black  powder,  insoluble 
in  water  but  soluble  in  acids,  with  which  it  forms  blue 
cupric  salts: 

CuO  +  H2S04  =  CuS04  +  H20. 

When  cupric  oxide  is  heated  alone,  it  does  not  readily 
yield  up  its  oxygen.     It  requires  the  high  temperature  of 
the  electric  arc  to  decompose  it  completely  into  oxygen  and 
copper.    But  when  heated  with  hydrogen  or  with  carbon,  its 
oxygen  is  easily  removed,  producing  water  or  carbon  dioxide : 
CuO  +  H2  =  Cu  +  H,0 
2CuO  +  C  =  2Cu  +  CO,. 

Cupric  oxide  is  much  employed  by  chemists  in  the  anal- 
ysis of  compounds  containing  carbon  and  hydrogen,  for  on 
being  heated  with  them  in  a  glass  tube  it  converts  their 
hydrogen  into  water  and  their  carbon  into  carbon  dioxide, 
both  of  which  are  easily  collected  separately  and  weighed. 
From  these  two  weights  the  percentages  of  hydrogen  and 
of  carbon  in  the  substance  heated  can  be  calculated. 

209.  Cuprous  compounds. — These  all  contain  the  linked 

Cu 
pair  of  copper  atoms,     |       They  are  almost  all  insoluble 

Cu 
in  water,  and  colorless. 

Cuprous  chloride,  Cu2Cl2,  can  be  made  by  shaking  a  solu- 
tion of  cupric  chloride,  CuCl2,  with  copper  filings.  It  is  a 
white  crystalline  powder  insoluble  in  water.  It  darkens 
on  exposure  to  light,  and  has  many  points  of  resemblance 
with  silver  chloride,  AgCl.  The  fact  that  the  density  of  its 
vapor  shows  that  the  formula  is  Cu2Cl2,  not  CuCl,  has 
been  noticed. 

Cuprous  bromide,  Cu2Br2,  and  cuprous  iodide,  Cu2I2,  are 
also  insoluble  in  water  and  decomposed  by  light.  They  are 
similar  to  silver  bromide  and  silver  iodide. 

210.  Cupric   compounds. — All   cupric   compounds — for 
example,  cupric  chloride,  CuCl2 — contain  one  bivalent  cop- 


THE  COPPER  GROUP  149 

per  atom.    When  they  are  dissolved  in  water  this  atom  be- 
comes an  ion  with  a  double  positive  charge.     Thus,  cupric 
chloride  in  dilute  water  solution  separates  into  three  ions — 
++    — 
Cu,  Cl,  and  Cl.    This  cupric  ion  communicates  a  blue  color 

to  all  solutions  of  cupric  salts,  and  the  color  of  dilute  water 
solutions  of  cupric  salts  is  the  same  no  matter  what  the  dis- 
solved salt  was,  not  only  the  same  on  simply  looking  at  the 
liquid,  but  the  same  also  when  light  which  has  passed 
through  the  liquid  is  sorted  out  into  its  colors  by  means  of 
a  prism.  This  is  an  interesting  proof  that  the  dissolved 
salts  have  separated  into  ions,  for  otherwise  different  salts 
would  give  different  colors  to  their  solutions.  This  remark 
applies  not  only  to  the  cupric  salts  but  to  all  metallic  salts 
whose  water  solutions  are  colored  (p.  131). 

211.  Different  series  of  salts  of  the  same  metal. — This  is 
the  first  time  we  have  had  an  opportunity  to  notice  a  re- 
markable fact — the  fact  that  many  metals  form  two  differ- 
ent series  of  salts,  the  two  sets  differing  greatly  in  their 
properties.  Solutions  of  the  cuprous  salts  are  just  as  un- 
like those  of  the  cupric  salts  in  color  and  in  chemical  con- 
duct as  though  they  contained  some  other  metal,  and  this 

+ 

Cu  ++ 

shows  us  that  the  cuprous  ion,     |     and  the  cupric  ion,  Cu, 

Cu 

are  very  different  things,  and  communicate  very  different 
properties  to  liquids  containing  them,  yet  both  are  nothing 
but  copper.  This  reminds  us  of  ordinary  oxygen  and  ozone, 
two  substances  composed  of  the  same  element  and  yet  very 
unlike. 

It  will  be  seen  that  in  a  cupric  ion,  Cu,  one  atom  of 
copper  carries  just  the  same  electric  charge  as  is  carried 

Cu 

by  two  atoms  in  the  cuprous  ion,     |        Therefore,  if  we 

Cu 


150  ELEMENTARY  CHEMISTRY 

allow  the  same  current  to  pass  through  a  cuprous  and  a 
cupric  solution,  the  amount  of  copper  precipitated  in  an 
hour  will  be  twice  as  great  in  the  euprous  solution  as  in 
the  cupric. 

212.  Reaction  of  cupric  compounds  with  metals. — When 
zinc  or  iron  is  placed  in  a  solution  of  a  salt  of  copper,  the 
metal  dissolves  and  copper  appears  in  its  place.     Thus,  if 
a  penknife-blade  is  dipped  into  a  solution  of  cupric  chloride, 
it  becomes  covered  at  once  with  red  metallic  copper.    This 
was  regarded  in  the  middle  ages  as  a  transformation  of  iron 
into  copper.    We  can  write  the  reaction — 

Fe  +  CuCl2  =  FeCl2  +  Cu. 
But  the  solution  of  cupric  chloride  consists  of  ions  Cu,  Cl, 

and  Cl,  and  it  is  clear  that  the  chlorine  ions  take  no  part  in 
the  process.  What  happens  is  simply  this.  An  iron  atom 
robs  a  copper  ion  of  its  charge  and  takes  its  place  in  the 
liquid,  the  copper  separating  as  metal,  thus: 

Fe  +  Cu  =  Fe  +  Cu. 

Since  the  liquid  already  contained  chlorine  ions,  and  now 
contains  iron  ions,  it  becomes  a  solution  of  iron  chloride, 
FeCl2. 

213.  Cupric  chloride,  CuCl2,  can  be  obtained  by  burning  cop- 
per in  chlorine,  or  by  dissolving  cupric  oxide  in  hydrochloric  acid, 
when  it  crystallizes  in  green  deliquescent  needles  of  the  composition 
CuClaSHaO,  which  are  very  soluble  in  water.     The  strong  solution 
of  cupric  chloride  is  grass-green,  because  the  salt  is  nearly  all  present 
as  molecules  CuCla,  not  separated  into  ions.     In  all  solutions  the 
amount  of  the  salt  which  is  so  separated  increases  rapidly  when  the 
liquid  is  diluted,  and  when  water  is  added  to  a  solution  of  cupric 

chloride  there  is  a  plain  change  from  the  bright  green  of  CuCla  to 

-"-+ 
the  blue  color  of  the  cupric  ion  Cu. 

Cupric  bromide,  CuBr2,  is  similar  to  the  chloride.  The  student 
should  notice  how  different  these  substances  are  from  the  corre- 
sponding cuprous  compounds. 


THE  COPPER  GROUP  151 

Cupric  iodide,  CuI2,  does  not  exist.  Various  reactions  by  which 
we  should  expect  it  to  be  formed  yield  instead  a  mixture  of  cuprous 
iodide  and  iodine : 

2CuI,  =  Cuja  +  I2. 

214.  Cupric  sulphate,  CuS04,  is  the  most  important  cu- 
pric  salt.  It  can  be  made  by  treating  copper  with  strong  sul- 
phuric acid  or  by  heating  copper  sulphide  gently  in  a  current 
of  air,  when  it  takes  up  oxygen.  It  is  a  white  powder,  which 
when  thrown  into  water  dissolves,  with  evolution  of  much 
heat,  to  a  blue  liquid.  When  this  solution  is  evaporated, 
blue  crystals  of  the  composition  CuS045H20  are  deposited. 
These  crystals  are  called  "  bluestone "  or  "  blue  vitriol." 
Large  quantities  of  bluestone  are  employed  in  copper-pla- 
ting and  in  filling  electric  batteries,  and  the  water  solution 
is  much  used  in  agriculture  for  the  destruction  of  parasites 
upon  plants. 


CHAPTER   XX 

THE  COPPER  GROUP  (Continued) 
SlLVEK,  Ag  =  108. 

215.  Occurrence  and  extraction, — Metallic  silver  occurs 
in  nature,  and  is  of  some  importance  as  an  ore.  The  chlo- 
ride AgCl — called  horn-silver — and  the  sulphide  Ag2S  are 
important  silver  ores.  Silver  is  very  commonly  associated 
with  lead,  and  most  lead  ores  contain  a  little  silver,  very 
often  enough  to  pay  for  its  extraction.  This  is  done  at  pres- 
ent by  melting  the  lead  obtained  from  such  ores  and  adding 
to  it  2  to  3  per  cent  of  zinc  and  stirring  thoroughly.  When 
the  vessel  is  allowed  to  stand  there  rises  to  the  surface  a 
layer  which  consists  of  an  alloy  of  zinc,  lead,  and  silver. 
This  is  removed,  and  the  lower  layer — mainly  lead — is 
treated  again  with  zinc  in  order  to  extract  the  last  traces 
of  silver. 

The  alloy  of  zinc,  lead,  and  silver  is  distilled,  when  the 
zinc  is  driven  off  as  vapor,  condensed,  and  used  again.  An 
alloy  of  lead  very  rich  in  silver  remains,  and  the  silver  is 
extracted  from  this  by  cupellation.  This  consists  in  melt- 
ing the  alloy  in  a  vessel  lined  with  bone-ash,  a  current  of 
air  being  drawn  over  the  surface.  The  lead  is  converted 
into  lead  oxide,  PbO,  which  is  partly  vaporized  and  part- 
ly absorbed  by  the  bone-ash.  When  the  amount  of  lead 
becomes  small,  brilliant  rainbow  hues  play  over  the  sur- 
face of  the  'melted  mass,  and  finally  the  bright  surface 
of  the  silver  beneath  appears.  Then  the  process  is 
stopped. 

152 


THE  COPPER  GROUP  153 

216.  Properties. — Silver  is  a  lustrous,  white  metal  which 
is  very  ductile  and  malleable,  and  a  better  conductor  of  heat 
and  the  electric  current  than  any  other  substance.  It  melts 
at  a  clear  yellow  heat  (1000°)  and  vaporizes  at  higher  tem- 
peratures. It  can  easily  be  distilled  in  the  oxyhydrogen 
flame,  in  lime  vessels,  and  this  method  has  been  employed 
in  the  preparation  of  the  pure  metal.  Stas,  when  he  needed 
pure  silver  to  determine  the  atomic  weight  of  the  metal, 
obtained  it  in  this  way.  He  informs  us  that  the  vapor  is 
pale  blue,  and  that  some  of  it  escaped  into  the  air,  making 
it  turbid  and  giving  it  a  metallic  taste. 

Silver  is  somewhat  sensitive  to  light,  and  when  a  glass 
plate  is  thinly  coated  with  it  and  exposed  in  a  camera  a 
picture  can  be  obtained.  This  shows  that  there  must  be  at 
least  two  allotropic  forms  of  the  metal,  one  of  which  is 
formed  by  the  action  of  light  on  the  other,  and,  in  fact,  a 
number  of  modifications  of  silver  have  been  described. 

Silver  is  not  affected  by  water  or  by  oxygen  at  any  tem- 
perature. We  have  seen  that  ozone  acts  upon  it,  convert- 
ing the  surface  into  black  Ag202.  A  silver  plate  exposed 
near  an  oxyhydrogen  flame  is  blackened  because  ozone  is 
formed  from  the  oxygen  by  the  high  temperature  and  then 
acts  upon  the  silver.  The  tarnishing  of  silver  vessels  in 
the  air  may  occasionally  be  due  to  the  action  of  ozone,  but 
it  is  usually  owing  to  the  production  of  a  film  of  silver 
sulphide,  Ag2S,  for  there  is  always  a  little  sulphur  in  coal, 
and  therefore  the  air  of  a  house  always  contains  traces  of 
gases  containing  sulphur,  which  come  from  the  burning  of 
coal  and  of  ordinary  gas,  which  is  made  from  it.  In  shops 
the  tarnish  is  prevented  by  keeping  the  silver  in  closed 
cases,  with  a  dish  containing  a  little  lime,  which  is  renewed 
from  time  to  time. 

Silver  is  scarcely  acted  upon  by  hydrochloric  or  dilute 
sulphuric  acid.  Strong  sulphuric  acid  converts  it  into  sil- 
ver sulphate,  Ag2S04,  and  nitric  acid  readily  dissolves  it  to 


154  ELEMENTARY  CHEMISTRY 

silver  nitrate,  AgN03.  Silver  may  be  regarded  as  univalent 
in  all  its  compounds. 

217.  Compounds  of  silver  with  oxygen. — Two  oxides  of 
silver  are  known  with  certainty. 

Silver  oxide,  Ag20,  is  a  brown  powder,  easily  decom- 
posed by  heat  into  silver  and  oxygen.  It  is  very  slightly  sol- 
uble in  water,  and  the  solution  has  an  unpleasant  metallic 

taste  and  distinct  basic  properties.    It  must  therefore  con- 

+ 
tain  hydroxyl  ions,  OH,  and  silver  ions,  Ag,  or,  what  is  the 

same  thing,  it  must  be  a  solution  of  silver  hydroxide,  AgOH. 
But  this  substance  has  never  been  obtained  solid.  Keactions 
by  which  we  should  expect  it  to  be  produced  always  yield  in- 
stead silver  oxide  and  water,  thus : 

2AgOH  =  Ag20  +  H20. 

Silver  dioxide,  Ag202,  which  we  may  write  Ag — 0 — 0 — Ag, 
has  been  obtained  by  the  action  of  ozone  on  silver  and  by 
other  methods.  It  crystallizes  in  black  needles,  which  are 
insoluble  in  water  and  easily  decomposed  by  heat. 

218.  Silver  sulphide,  Ag2S,  can  be  obtained  in  blackish- 
gray  crystals  by  heating  finely  divided  silver  in  the  vapor 
of  sulphur.     Oxidized  silver  is  a  misleading  name  given  to 
silver  which  has  been  thinly  coated  with  silver  sulphide  by 
boiling  it  in  a  solution  containing  potassium  pentasulphide, 
K2S6. 

219.  Compounds  of  silver  with  fluorine,  chlorine,  bro- 
mine, and  iodine. — These  four  elements  closely  resemble 
each  other  chemically,  and  it  is  convenient  to  consider  their 
silver  compounds  together. 

Silver  fluoride,  AgF,  differs  remarkably  from  the  other  three 
compounds.  It  is  a  clear  yellow,  horny  mass,  which  turns  black  in 
the  light,  is  deliquescent,  and  very  soluble  in  water.  When  silver 
is  suspended  in  this  solution  a  beautifully  crystallized  salt,  called 
silver  subfluoride,  Ag3F,  is  produced. 

It  has  been  pointed  out  that  complete  insolubility  is  a  state  of 


JOHN   D ALTON 
B.  England,  1766.     D.  1844. 


THE  COPPER  GROUP  155 

things  which  real  substances  approach  but  probably  never  reach. 
Silver  chloride,  bromide,  and  iodide  are  interesting  examples  of  this 
fact.  It  takes  something  like  a  million  parts  of  water  to  dissolve 
one  part  of  silver  chloride,  and  the  others  are  still  less  soluble,  yet  it 
has  been  shown  that  water  dissolves  small  quantities  of  all  of  them, 
and  that  the  solubility  decreases  in  the  order  in  which  the  com- 
pounds have  been  named. 

220.  Silver  chloride,  AgCl,  separates  as  a  heavy,  white, 
curdy  precipitate  when  a  liquid  containing  chlorine  ions  is 
mixed  with  one  containing  silver  ions : 

A+g  +  01  =  AgCl. 

Thus,  when  a  solution  of  table  salt  is  mixed  with  a  solution 
of  silver  nitrate : 

NaCl  +  AgN03  =  AgCl  +  NaN03. 

Silver  chloride  turns  bluish-gray,  and  finally  black,  when 
exposed  to  light.  This  is  important  in  photography,  and 
has  been  known  for  more  than  a  century,  yet  we  are  still 
uncertain  with  respect  to  what  happens  when  the  darkening 
occurs. 

Silver  bromide,  AgBr,  and  silver  iodide,  Agl,  resemble 
silver  chloride.  They  are  yellow,  and  like  the  chloride, 
darken  when  exposed  to  light. 

PHOTOGRAPHY 

221.  The  daguerreotype  process. — The  sensitiveness  of 
these  three  silver  compounds  to  light  is  the  basis  of  the 
photographic  process.     The  first  attempt  to  apply  the  fact 
to  picture-making  which  attained  any  practical  importance 
was  that  of  Daguerre.    He  coated  a  silver  plate  with  a  film 
of  silver  iodide  by  allowing  the  vapor  of  iodine  to  act  upon 
it,  and  then  exposed  the  sensitive  surface  so  obtained  in  a 
camera ;  but  the  exposure  required  was  long  and  the  results 
unsatisfactory  until  he  discovered,  quite  by  accident,  that 
if,  after  the  exposure,  the  plate  was  treated  with  the  vapor 


156  ELEMENTARY   CHEMISTRY 

of  mercury,  the  mercury  would  condense  upon  those  parts 
of  the  surface  which  had  been  acted  upon  by  the  light,  and 
the  picture  appear.  Then  the  plate  was  treated  with  a 
liquid  which  would  dissolve  away  the  unaltered  silver  iodide 
from  those  portions  upon  which  the  light  had  not  acted. 
This  was  necessary,  because  otherwise  the  plate  would  have 
blackened  uniformly  on  exposure  to  light,  destroying  the 
picture. 

222.  The  modern  dry  plate  consists  of  a  plate  of  glass, 
one  side  of  which  is  covered  with  a  film  of  hardened  gelatin 
containing  finely  divided  silver  bromide,  AgBr.  This  is 
placed  in  the  camera  in  such  a  way  that  this  surface  is  in 
the  focus  of  the  lens,  and  receives  the  image.  After  the 
exposure  to  light  the  plate  is  just  the  same  in  appearance 
as  before.  It  is  only  on  being  treated  with  a  liquid  called 
the  "  developer "  that  the  picture  makes  its  appearance. 
A  developer  is  a  substance  which  has  this  peculiar  property, 
that  while  it  has  no  action  upon  silver  bromide  which  has 
been  kept  in  the  dark,  it  converts  silver  bromide  which  has 
been  exposed  to  light  into  metallic  silver.  Accordingly, 
silver  bromide  which  has  been  exposed  to  light  must  have 
sustained  some  alteration  which  causes  it  to  be  attacked 
by  the  developer  afterward.  In  spite  of  much  careful  work 
upon  the  problem,  the  nature  of  this  change  is  still  disputed. 
This  much  is  known,  that  bromine  is  set  free  when  light  acts 
upon  silver  bromide.  Perhaps  the  light  decomposes  very 
small  quantities  of  the  silver  bromide  into  silver  and  bro- 
mine, leaving  traces  of  finely  divided  silver  wherever  it  acts 
upon  the  plate.  Then,  in  the  developing  process,  the  silver 
produced  by  the  action  of  the  developer  on  the  silver  bro- 
mide of  the  plate  crystallizes  upon  these  silver  particles,  just 
as  a  supersaturated  solution  crystallizes  upon  a  particle  of 
the  same  substance  thrown  into  it  (p.  14).  If  this  view  is 
correct,  the  quantity  of  silver  produced  must  be  very  small, 
for  when  a  plate  is  exposed  to  light,  and  then  the  silver 


THE  COPPER  GROUP 


157 


bromide  dissolved  out  of  it  without  any  developing,  no  silver 
can  be  detected  even  under  the  microscope. 

Or  perhaps  the  light  converts  the  silver  bromide  into 
silver  sub-bromide,  Ag2Br  (analogous  to  silver  sub-fluoride, 
Ag2F),  which  is  then  converted  into  silver  by  the  developer. 

223.  At  any  rate,  wherever  the  light  has  fallen  upon  the 
plate,  the  developer  produces  dark  metallic  silver.  The 
next  step  is  to  dissolve  away  the  unaltered  silver  bromide 
in  order  to  prevent  it  from  being  acted  upon  by  light  and 
spoiling  the  image.  This  is  called  "  fixing,"  and  is  accom- 


FIQ.  34,  a.— Positive.  FIG.  34,  ft.— Negative. 

plished  by  soaking  the  plate  in  a  solution  of  sodium  thio- 
sulphate,  N~a2S203,  called  "  hypo "  by  the  photographer. 
The  plate  is  now  called  a  "  negative,"  for  the  lights  and  shad- 
ows are  exactly  reversed  in  it.  The  bright  parts  of  the  image 
are  covered  with  blackish  silver,  while  the  dark  portions  are 
simply  clear  glass  (Fig.  34,  &). 


158  ELEMENTARY  CHEMISTRY 

224.  After  being  thoroughly  washed  and  dried  the  nega- 
tive is  ready  for  " printing"  It  is  exposed  to  light,  with 
a  piece  of  sensitive  paper  back  of  it,  in  contact  with  the 
image.  The  light  passes  easily  through  the  clear  portions, 
which  therefore  become  dark  in  the  print,  but  is  arrested 
by  those  parts  which  are  covered  with  opaque  silver,  which 
therefore  remain  white  on  the  paper.  Hence  the  lights  and 
shadows  are  again  reversed,  and  are  now  the  same  as  they 
were  in  the  original  scene. 

Various  papers  are  employed' for  printing.  A  very  com- 
mon one  consists  of  paper  covered  with  a  film  of  albumin 
or  gelatin  containing  silver  chloride,  AgCl.  We  have  seen 
that  silver  chloride  darkens  when  it  is  exposed  to  light. 
There  is  still  doubt  as  to  what  happens  to  it  when  it  is 
alone,  but  in  presence  of  the  albumin  silver  is  produced. 
Silver  obtained  in  this  way  has  an  unpleasant  reddish  color, 
so  the  paper  is  afterward  dipped  in  a  solution  containing 
gold  chloride.  In  this  way  part  of  the  silver  of  the  image 
is  replaced  by  gold,  which  gives  a  more  satisfactory  color. 
This  is  called  toning. 

The  "  developing  "  papers,  like  "  Velox,"  work  on  the 
same  principle  as  a  plate.  The  paper  is  covered  with  a 
film  of  gelatin  containing  silver  bromide,  and  when  it  is 
exposed  back  of  a  negative  the  paper  remains  white,  but 
the  picture  appears  upon  dipping  the  paper  into  the  de- 
veloper. 

Both  these  kinds  of  paper  require  to  be  "  fixed  "  after 
the  picture  appears.  In  other  words,  the  unaltered  silver 
chloride,  or  silver  bromide,  as  the  case  may  be,  must  be  re- 
moved, since  otherwise  the  whole  paper  would  shortly  turn 
black.  This  is  accomplished  in  just  the  same  way  as  with 
the  plate,  by  soaking  the  paper  in  a  solution  of  sodium 
thiosulphate  and  washing  thoroughly. 


CHAPTER   XXI 

THE  COPPER   GROUP  (Continued) 

GOLD,  Au  =  197. 

225.  Occurrence  and  extraction. — Gold  occurs  free  in 
nature,  and  was  therefore  among  the  earliest  metals  known 
to  man.  It  is  rather  widely  distributed,  but  is  found  abun- 
dantly in  only  a  few  localities.  The  great  gold  deposits 
are  in  South  Africa,  Australia,  the  United  States,  and  Rus- 
sia. Under  ordinary  circumstances  South  Africa  stands 
first  as  regards  the  amount  of  gold  produced,  though  of 
late  the  interruption  of  mining  by  war  in  the  Transvaal  has 
allowed  the  United  States  to  take  first  rank. 

Gold  occurs  chiefly  in  two  ways,  either  in  grains  or 
fragments  scattered  through  beds  of  gravel  or  sand  (placer 
deposits),  or  else  disseminated  through  some  compact  rock, 
usually  quartz.  In  working  a  placer  deposit  on  a  large 
scale  a  powerful  jet  of  water  is  thrown  into  the  mass.  A 
torrent  of  water  charged  with  earth  and  stones  rushes  out, 
and  is  led  away  through  a  long  inclined  channel  built  of 
boards.  Pieces  of  wood  projecting  above  the  general  level 
are  nailed  transversely  across  the  bottom  of  this  channel 
at  intervals.  The  gold,  being  far  denser  than  the  rest  of 
the  solid  matter  carried  by  the  water,  sinks  most  rapidly 
and  collects  back  of  these  projections.  From  time  to  time 
the  water  is  stopped  and  the  gold  removed. 

A  compact  rock  containing  gold  is  crushed  finely  and 
then  brought  into  contact  with  mercury.  Gold  forms  an 
alloy — an  amalgam,  as  it  is  called — with  mercury  more 

159 


160  ELEMENTARY   CHEMISTRY 

readily  than  any  other  metal.  This  amalgam  is  removed 
and  distilled.  The  mercury  vaporizes  and  is  condensed, 
while  the  gold  remains  and  is  further  purified. 

226.  Properties. — Gold  is  a  metal  with  a  characteristic 
yellow  color   and,   when   polished,   a   brilliant   luster.      It 
has  a  very  low  specific  heat,  and  therefore  feels  warmer 
to  the  hand  than  copper  or  iron  at  the  same  tempera- 
ture.    It  is  marvelously  malleable  and  ductile.     Gold  is 
harder   than   tin   but   softer    than    silver,    and    an    object 
made  of  it  would  wear  away  rapidly  in  use.     Therefore 
the  gold  for  use  in  coinage  is  alloyed  with  10  per  cent 
of  copper,  and  jeweler's  gold  with  copper  or  copper  and 
silver. 

When  heated,  gold  melts  at  a  somewhat  higher  tempera- 
ture than  silver,  and  at  the  temperature  of  the  electric  arc 
vaporizes  rapidly,  torrents  of  greenish-yellow  smoke  emerg- 
ing from  the  vessel.  It  is  not  affected  by  oxygen,  by  water, 
or  by  air  at  any  temperature,  nor  is  it  attacked  by  the 
ordinary  acids.  It  is  said  that  finely  divided  gold  dissolves 
in  very  strong  nitric  acid,  and  is  precipitated  again  when 
water  is  added.  Although  compact  gold  is  not  affected 
by  either  hydrochloric  or  nitric  acids  it  dissolves  readily  in 
the  mixture  of  the  two.  This  mixture  is  called  "  aqua 
regia,"  and  it  also  dissolves  platinum  and  some  other  met- 
als which  are  not  attacked  by  either  acid  alone.  This  is 
because  the  nitric  acid  liberates  chlorine  from  the  hydro- 
chloric acid,  and  the  chlorine  converts  the  metal  into  its 
chloride. 

227.  Uses. — On  account  of  its  appearance  and  its  resist- 
ance to  the  atmosphere,  gold  is  extensively  used  for  plating 
other  metals.     The  object  to  be  gilded  is  thinly  covered 
with  gold  amalgam  *  and  then  carefully  heated  to  drive  off 
the  mercury.    Or  the  object  is  suspended  in  a  solution  con- 

1  The  alloys  of  mercury  with  other  metals  are  called  amalgams. 


THE  COPPER  GROUP  161 

taining  gold  and  made  the  negative  pole  of  a  feeble  cur- 
rent passing  through  the  liquid.  The  first  method  is  called 
"  fire-gilding "  and  the  second  "electro-plating,"  and  by 
either  an  exceedingly  thin  film  of  gold  can  be  deposited. 

228.  Gold  can  be  either  univalent  or  trivalent.     The  compounds 
in  which  it  is  univalent — the  aurous  compounds — are  much  like  the 
compounds  of  silver.     Thus,  aurous  chloride,  AuCl,  aurous  bromide, 
AuBr,  and  aurous  iodide,  Aul,  are  white  or  yellow,  and  insoluble  in 
water.      They  are  easily  decomposed  ly  heat,  leaving  a  residue  of  gold, 
and  this,  in  fact,  is  true  of  all  gold  compounds  without  exception. 

229.  The  compounds  in  which  gold  is  trivalent,  the 
auric  compounds,  are  the  most  important. 

Auric  oxide,  Au203,  is  a  blackish-brown  powder,  which 
at  250°  separates  into  gold  and  oxygen.1  When  ammonia 
water  is  poured  over  it,  it  is  converted  into  "fulminating 
gold"  a  green  powder  which  explodes  most  violently  on 
being  heated  or  struck. 

Auric  hydroxide,  Au(OH)3,  is  a  yellow  powder  insoluble 
in  water.  At  100°  water  separates  from  it  and  auric  oxide 
remains : 

2Au(OH)3  =  Au203  +  3H20. 

Auric  hydroxide  behaves  more  like  an  acid  than  a  base. 
Its  hydrogen  can  be  replaced  by  metals  like  sodium  and 
potassium,  forming  salts  called  aurates.  This  shows  that, 
while  gold  is  a  perfect  metal  physically,  its  metallic  char- 
acter chemically  is  not  well  marked. 


1  In  statements  like  this  it  is  understood  that  the  substance  is 
heated  in  the  air.  The  temperature  at  which  auric  oxide  will  yield  up 
its  oxygen  depends  entirely  upon  the  atmosphere  in  which  it  is  heated. 
Thus,  if  sealed  up  in  a  strong  glass  tube  with  compressed  oxygen,  a 
higher  temperature  than  250°  would  be  needed  to  decompose  it,  while, 
on  the  other  hand,  when  it  is  heated  in  a  vacuum,  the  temperature 
required  is  not  so  high  as  in  the  air.  The  more  free  oxygen  in  the 
gas  in  contact  with  it  the  higher  the  temperature  required. 


162  ELEMENTARY  CHEMISTRY 

Auric  chloride,  AuCl3,  is  made  by  dissolving  gold  in 
aqua  regia,  or  better,  by  treating  finely  divided  gold  with 
chlorine  and  dissolving  the  product  in  water,  which  is  then 
evaporated.  Auric  chloride  forms  brown,  imperfect  crystals 
freely  soluble  in  water.  Heat  separates  it  into  gold  and 
chlorine. 


CHAPTEK   XXII 

THE  CALCIUM  GROUP 
Calcium,  Ca.  Strontium,  Sr.  Barium,  Ba. 

230.  General  properties  of  the  calcium  group. — Calcium, 
strontium,,  and  barium  are  strongly  metallic.    They  oxidize 
rapidly  in  the  air,  and  decompose  water  energetically,  pro- 
ducing their  hydroxides  and  liberating  hydrogen,  thus : 

Ca  +  2H20  =  Ca(OH)2  +  H2. 

The  hydroxides  are  very  strong  bases,  barium  hydroxide 
being  the  strongest. 

It  will  be  noticed  that  there  is  a  considerable  similarity 
between  these  metals  and  the  metals  of  the  sodium  group. 
The  most  important  difference  is  that  these  metals  are  al- 
most always  bivalent,  and  this  will  enable  us  to  write  the 
formulas  of  most  of  their  compounds  without  memorizing 
them. 

CALCIUM,  Ca  =  40. 

231.  Occurrence. — Calcium  stands  fifth  among  the  ele- 
ments  in  point   of  abundance,  making  up   nearly  4   per 
cent  of  the  earth's  crust.    Limestone  and  chalk  are  calcium 
carbonate,  CaC03,  and  marble  is  the  same  material  in  com- 
pact crystalline  state.    The  pearl  is  also  calcium  carbonate, 
and  owes  its  value  to  its  beautiful  luster,  not  to  the  material 
of  which  it  consists.     Most  shells — those  of  oysters  and 
clams,  for  instance — consist  of  calcium  carbonate  with  more 
or  less  organic  matter.1     Compounds  of  calcium  are  con- 

1  By  organic  matter  we  mean  complicated  compounds  of  carbon, 
hydrogen,  nitrogen,  and  oxygen,  originating  in  organized  beings — 
12  163 


164  ELEMENTARY  CHEMISTRY 

tained  in  most  common  rocks,  in  all  soils,  and  in  the  bodies 
of  most  plants  and  animals.  In  the  latter  it  exists  chiefly 
in  the  bones  and  teeth,  which  contain  about  half  their  weight 
of  calcium  phosphate,  Ca3(P04)2,  the  rest  being  mainly 
organic  matter.  When  bones  are  burned,  the  organic 
matter  is  consumed  and  the  calcium  phosphate  is  left  as 
an  ash. 

232.  Preparation  and  properties. — Compounds  of  cal- 
cium, like  calcium  carbonate  and  lime — which  is  calcium 
oxide,  CaO — have  been  known  for  centuries,  but  the  metal 
has  only  been  obtained  pure  and  in  large  quantities  quite 
recently  (1898).  This  was  accomplished  by  heating  cal- 
cium iodide,  CaI2,  to  redness  with  sodium  in  a  crucible : 

CaI2  +  2Na  =  2NaI  +  Ca. 

When  the  mass  cools  it  is  powdered  and  treated  with  pure 
alcohol,  which  dissolves  the  sodium  iodide,  leaving  the  cal- 
cium as  a  gray  powder,  which  can  be  made  into  a  compact 
mass  by  melting  it  in  a  vacuum. 

Calcium  is  a  silver-white  metal  with  a  brilliant  luster, 
which  it  rapidly  loses  in  moist  air.  It  is  somewhat  harder 
than  lead,  and  less  than  twice  as  dense  as  water.  Heated 
in  a  vacuum,  it  melts  at  a  red  heat.  Heated  in  oxygen, 
it  takes  fire  and  burns  with  blinding  brilliancy  to  lime — 
calcium  oxide,  CaO — and  the  temperature  produced  is  so 
high  that  the  lime  melts.  When  calcium  is  burned  in  the 
air  it  combines  not  only  with  the  oxygen  but  also  with 
the  nitrogen,  and  the  product  is  a  mixture  of  lime  with 
calcium  nitride,  Ca3N2.  When  calcium  is  thrown  into  water 
hydrogen  quickly  escapes  and  a  white  mass  of  calcium  hy- 
droxide, Ca(OH)2 — commonly  called  "slaked  lime" — is 
formed. 


animals  or  plants.     Albumin  and  gelatin  are  examples.     The  chem- 
ical nature  of  such  substances  is  not  yet  fully  understood. 


THE  CALCIUM  GROUP  165 

233.  Calcium  oxide,  CaO  (lime),  is  the  product  of  the 
burning  of  calcium  in  oxygen.  Lime  is  made  on  the  large 
scale  by  heating  some  form  of  calcium  carbonate  in  such  a 
way  that  the  carbon  dioxide  produced  can  readily  escape: 

CaC03  =  CaO  +  C02. 

Oyster-shells,  limestone,  or  marble  is  mixed  with  coal 
and  thrown  in  at  the  top  of  a  furnace — the  lime-kiln — the 
lime  produced  being  removed  at  the  bottom  and  new  mate- 
rials supplied  above  so  that  the  process  is  continuous.  Cases 
of  suffocation  sometimes  occur  during  the  winter  among 
people  who,  attracted  by  the  warmth,  attempt  to  sleep  near 
a  lime-kiln,  and  the  equation  above  helps  us  to  understand 
this,  for  carbon  dioxide,  C02,  is  a  suffocating  gas. 

If  calcium  carbonate,  CaCO8,  is  heated  in  a  closed  vessel  which  it 
completely  fills,  it  can  be  melted  without  this  decomposition  occur- 
ring. If  the  calcium  carbonate  only  partly  fills  the  vessel  in  which 
it  is  heated,  some  lime  will  be  produced,  and  carbon  dioxide  will 
accumulate  in  the  space  above  the  solid.  The  bottom  of  the  vessel 
will  contain  a  solid  mixture  of  lime  and  calcium  carbonate,  and 
above  this  there  will  be  a  space  containing  gaseous  carbon  dioxide. 
The  weight  of  carbon  dioxide  in  each  cubic  centimeter  of  this  space, 
or,  what  is  exactly  the  same  thing%  the  pressure  of  the  gas  against 
the  walls  of  the  vessel,  depends  entirely  upon  the  temperature,  and 
is  not  in  the  least  affected  by  the  size  of  the  vessel  nor  by  the  quan- 
tity of  lime  or  calcium  carbonate  in  it.  For  every  temperature  the 
pressure  of  the  carbon  dioxide  against  the  vessel  walls — or  the  quan- 
tity of  the  gas  in  each  cubic  centimeter — has  a  certain  value,  and 
when  that  value  has  once  been  reached  years  of  heating  to  that  tem- 
perature will  not  change  it  nor  decompose  any  more  calcium  car- 
bonate. The  higher  the  temperature  the  more  calcium  carbonate  is 
decomposed  and,  of  course,  the  greater  the  pressure.  But  when 
calcium  carbonate  is  heated  in  an  open  vessel,  carbon  dioxide  con- 
stantly escapes,  and  the  decomposition  becomes  complete.  This 
reminds  us  of  the  behavior  of  water  evaporating  into  an  enclosed 
space ;  we  have  seen  that  water  evaporates  until  the  pressure  of  the 
water-vapor  reaches  a  value  which  is  always  the  same  at  the  same 
temperature.  This  is  a  physical  change,  while  the  decomposition 


166  ELEMENTARY  CHEMISTRY 

of  calcium  carbonate  by  heat  is  a  chemical  change,  yet  the  two 
processes  have  much  in  common. 

Common  lime  is  gray,  but  the  pure  substance  is  white. 
When  treated  with  the  oxyhydrogen  flame,  it  glows  brightly 
— the  lime-light — but  is  not  affected  otherwise.  At  the 
temperature  of  the  electric  arc  it  can  readily  be  melted  and 
boiled.  When  liquid  lime  is  allowed  to  cool,  it  solidifies  to 
a  milky  crystalline  mass,  and  when  a  piece  of  lime  is  heated 
for  a  time,  not  quite  hot  enough  to  melt  it,  it  becomes 
covered  with  crystals,  which  take  the  form  of  cubes.  Lime, 
when  exposed  to  the  air,  slowly  absorbs  water  and  carbon 
dioxide,  and  falls  to  a  white  powder,  called  "  air-slaked 
lime,"  which  is  a  mixture  of  calcium  hydroxide  and  calcium 
carbonate : 

CaO  +  H20  =  Ca(OH)2 

CaO  +  C02  =  CaC03. 

Another  compound  of  calcium  and  oxygen  is  known.  This  is 
calcium  dioxide,  CaOa,  which  is  white  and  easily  separated  into  lime 
and  oxygen  by  heat.  If  we  wish  to  consider  caloium  as  always 
bivalent — it  is  certainly  almost  always  so — we  can  write  this  sub- 
stance thus: 


234.  Calcium  hydroxide,  Ca(OH)2,  "slaked  lime,"  is 
made  by  the  union  of  lime  with  water.  Much  heat  is  pro- 
duced, and  this  may  be  shown  by  placing  on  the  surface  of 
the  lime  a  little  gunpowder,  which,  if  care  is  taken  not 
to  wet  it,  is  inflamed.  Calcium  hydroxide  is  a  loose  white 
powder,  about  twice  as  dense  as  water,  in  which  it  is  only 
slightly  soluble.  It  is  one  of  the  very  few  solids  which  are 
more  soluble  in  cold  water  than  in  hot.  The  saturated 
water  solution  of  calcium  hydroxide  is  called  lime-water,  and 
is  used  in  medicine.  On  account  of  the  small  solubility  of 
calcium  hydroxide,  lime-water  contains  very  little  of  it,  only 
about  one  part  in  seven  hundred. 


THE  CALCIUM  GROUP  167 

Calcium  hydroxide  slowly  absorbs  carbon  dioxide  from 
the  air,  passing  into  calcium  carbonate : 

Ca(OH)2  +  C02  =  CaC03  +  H20. 

The  use  of  slaked  lime  for  building  depends  upon  this 
fact.  The  first  hardening  of  mortar — "  setting  " — is  sim- 
ple drying.  Then  follows  a  slow  absorption  of  carbon  di- 
oxide from  the  air,  forming  a  cement  which  finally  binds 
the  grains  of  sand  together  in  a  mass  of  stony  hardness. 
This  change  of  calcium  hydroxide  into  calcium  carbonate 
begins  at  the  surface  and  proceeds  very  gradually  toward 
the  interior,  for  the  examination  of  ancient  samples  of 
mortar — six  centuries  old  or  more — has  shown  that  calcium 
hydroxide  is  still  present  in  the  inner  portions. 

Lime-water  is  alkaline  to  litmus  and  other  indicators, 
and  must  therefore  contain  hydroxyl  ions.  Calcium  hy- 
droxide— like  all  the  hydroxides  in  this  group — is  a  strong 
base,  reacting  with  acids  to  produce  calcium  salts  and  water, 
thus: 

Ca(OH)2  +  2HC1     =  CaCl2  +  2H20, 
Ca(OH)2  +  2HN03  =  Ca(N03)2  +  2H.0.1 

When  calcium  hydroxide  is  heated,  steam  is  given  off  and 
lime  is  left: 

Ca(OH)2  =  CaO  +  H20. 

235.  Compounds  of  calcium  containing  sulphur. — Cal- 
cium sulphide,  CaS,  is  produced  when  the  metal  is  heated  in 
sulphur  vapor,  in  which  it  burns  brilliantly.  It  is  made  by 
heating  a  mixture  of  calcium  sulphate  and  charcoal : 

CaS04  +  2C  =  CaS  +  2C02. 

It  is  a  white  solid,  very  slightly  soluble  in  water.    When 
impure,  calcium  sulphide  possesses  the  property  of  being 

1  The  writing  of  equations  like  these  will  become  a  simple  matter 
for  the  student  if  he  will  remember  that  calcium  is  bivalent,  and, 
accordingly,  one  Ca  takes  the  place  of  two  H. 


168  ELEMENTARY  CHEMISTRY 

luminous  in  the  dark  after  light  has  acted  upon  it.  It  is 
the  basis  of  phosphorescent  paint. 

Calcium  sulphate,  CaS04,  occurs  in  nature  as  the  mineral 
anhydrite.  The  very  common  mineral  gypsum  is  calcium 
sulphate  with  two  equivalents  of  water  of  crystallization — 
CaS04,  2H20.  When  gypsum  is  gently  heated  it  gives  off 
part  of  this  water  and  falls  to  a  white  powder,  which  is 
called  "  plaster  of  Paris."  If  this  powder  is  wet  it  again 
combines  with  the  water  and  sets  to  a  hard  mass.  In  this 
way  plaster  casts  are  made.  But  if  the  gypsum  is  too  highly 
heated  in  the  first  place,  it  will  not  combine  again  with 
water,  or  only  very  slowly,  and  is  spoiled  for  this  purpose. 
It  is  then  said  to  be  "  dead-burnt." 

Calcium  sulphate  is  not  very  soluble  in  water,  and  sepa- 
rates as  a  white  precipitate  when  any  liquid  containing  S04 

ions — for  instance,  sulphuric  acid,  or  a  water  solution  of 

++ 
any  sulphate — is  added  to  a  solution  containing  Ca  ions  (a 

solution  of  a  calcium  salt) .  When  sulphuric  acid  is  added  to 
a  strong  solution  of  calcium  chloride  the  liquid  solidifies  to 
a  white  mass: 

CaCl2  +  H2S04  =  CaS04  +  2HC1, 
torrents  of  hydrochloric  acid  gas  being  given  off. 

236.  Calcium  fluoride,  CaF2,  is  the  common  mineral 
fluorite  usually  called  fluorspar.  It  crystallizes  in  cubes, 
and  is  transparent  and  colorless  when  pure,  but,  like  many 
other  minerals,  it  is  often  colored  blue,  yellow,  or  some  other 
color  by  traces  of  impurities  about  whose  nature  there  is 
still  dispute.  When  heated  in  the  dark,  the  crystals  become 
luminous. 

Calcium  fluoride  is  almost  insoluble  in  water.  When 
heated  with  strong  sulphuric  acid,  it  yields  calcium  sulphate 
and  hydrofluoric  acid,  HF,  and  it  is  employed  in  the  pro- 
duction of  this  acid — 

CaF2  +  H2S04  =  CaS04  +  2HF. 


THE  CALCIUM  GROUP  169 

237.  Calcium    chloride,    CaCl2,    results    when    calcium 
burns  in  chlorine.    It  is  made  by  dissolving  marble  in  hydro- 
chloric acid : 

CaC03  +  2HC1  =  CaCl2  +  H20  +  C02. 

On  evaporating  and  cooling  the  liquid  it  separates  in 
colorless  crystals,  which  contain  six  molecular  weights  of 
water,  CaCl26H20.  When  these  are  heated,  water  escapes 
and  anhydrous  calcium  chloride  remains  as  a  white  mass. 
This  absorbs  water  energetically,  and  it  is  much  used  in 
the  laboratory  for  drying  gases  and  liquids.  It  is  very 
soluble  in  water,  which  dissolves  at  100°,  about  1-J 
times  its  own  weight,  and  much  more  at  higher  tempera- 
tures. 

238.  Calcium  carbonate,  CaC03,  is  the  most  abundant 
compound  of  calcium.     Compact,  not  crystalline,  and  im- 
pure, it  is  limestone,  one  of  the  most  common  rocks.    From 
this — probably  by  the  combined  action  of  great  pressure, 
moderately  high  temperature  and  water — is  formed  marble, 
which  is  a  denser,  purer  variety,  composed  of  a  mass  of 
crystals  so  crowded  together  that  none  has  had  a  chance 
to  develop  properly.     Calcium  carbonate  in  the  shape  of 
minute  shells  pressed  together,  forming  a  soft  white  rock, 
is  called  chalk.    When  a  grain  of  sand,  or  some  other  small 
hard  particle,  gets  inside  the  shell  of  an  oyster,  there  is 
deposited  around  it  layer  after  layer  of  calcium  carbonate, 
and  so  is  formed  the  pearl,  which  owes  its  luster  to  this 
structure  in  layers.     The  shells  of  shell-fish  consist  of  cal- 
cium carbonate,  and  the  same  is  true  of  the  coral  rock  of 
which  coral  islands  are  built  up.     Stalactites  and  stalag- 
mites have  the  same  composition.     Calcium  carbonate,  al- 
though insoluble  in  water,  dissolves  in  water  containing 
carbon  dioxide.    When  such  a  solution  oozes  from  the  rocks 
forming  the  roof  of  a  cavern,  the  carbon  dioxide  escapes 
into  the  air  and  the  calcium  carbonate,  which  thereupon 


170  ELEMENTARY  CHEMISTRY 

deposits,  forms  the  stalactite,  while,  at  the  place  where 
the  drip  strikes  the  floor,  the  stalagmite  grows  up  to 
meet  it. 

Calcium  compounds  which  vaporize  in  the  Bunsen  flame 
communicate  to  it  a  strong  orange  color.  Compounds  like 
lime,  which  do  not  volatilize  in  the  flame,  do  not  color  it,  but 
the  color  can  be  obtained  by  first  moistening  the  substance 
with  hydrochloric  acid.  Calcium  chloride  is  thus  produced, 
and  vaporizes  when  placed  in  the  flame,  giving  the  orange 
color. 

239.  Strontium  and  barium. — These  metals  are  similar 
to  calcium,  but  are  harder  and  denser.  Strontium  is  yellow- 
ish and  barium  silver  white.  Both  burn  brilliantly  to  their 
oxides  when  heated  in  the  air,  and  both  are  rapidly  con- 
verted into  their  hydroxides  by  water,  thus : 

Ba  +  2H20  =  Ba(OH)2  +  H2. 

Strontium  compounds  color  the  burner-flame  red  and 
barium  compounds  green.  Red  fire  is  a  mixture  of  stron- 
tium nitrate,  Sr(N03)2,  with  some  combustible  substance 
like  powdered  sulphur,  charcoal,  or  shellac.  The  materials 
are  powdered  separately  and  carefully  mixed.  When  the 
mixture  is  lighted,  the  combustible  material  burns  at  the 
expense  of  the  oxygen  of  the  nitrate,  and  the  strontium  com- 
pounds  which  are  carried  into  the  flame  produce  the  red 
color.  Green  fire  is  a  similar  mixture  made  with  barium 
nitrate  instead  of  the  strontium  salt. 

Barium  sulphate,  BaS04,  is  almost  entirely  insoluble  in 
water,  and  is  formed  as  a  dense  white  crystalline  precipi- 

++ 
tate  when  Ba  ions  and  S04  ions  are  introduced  into  the 

same  liquid — for  instance,  when  a  solution  of  barium  chlo- 
ride is  mixed  with  dilute  sulphuric  acid : 

BaCl2  +  H2S04  =  BaS04  +  2HC1,  or 


THE  CALCIUM  GROUP 

This  reaction  is  much  employed  as  a  test  for  barium,  and 
still  more  as  a  test  for  sulphuric  acid  and  sulphates — that  is, 

for  S04  ions.  Barium  sulphate  is  insoluble  in  acids  and  in 
ammonia,  and  this  aids  us  to  distinguish  it  from  other  pre- 
cipitates of  similar  appearance. 

Radium  is  a  recently  discovered  metal  belonging  to  the  calcium 
group.  Its  compounds  are  rare  in  nature.  They  are  found  chiefly 
as  an  impurity  in  pitchblende  (page  232),  and  even  this  mineral  con- 
tains at  most  one  part  in  a  million  of  radium. 

Compounds  of  radium  Hre  luminous  in  the  dark,  and  the  less 
pure  specimens  appear  to  retain  their  light-giving  power  for  years 
without  loss.  They  continuously  give  off  heat,  and  if  a  tube  con- 
taining a  radium  compound  is  packed  in  wool,  so  as  to  prevent  loss 
of  heat,  it  will  maintain  itself  at  a  temperature  several  degrees 
higher  than  that  of  surrounding  objects.  Radium  compounds  con- 
stantly send  out  rays  which  travel  in  straight  lines,  and  which  affect 
the  photographic  plate  in  somewhat  the  same  way  as  light.  How- 
ever, these  rays. pass  easily  through  substances  which  are  quite 
opaque  to  light,  like  black  paper,  aluminium,  and  copper.  They 
turn  glass  violet  or  black,  and  cause  dry  air  and  other  non-con- 
ducting substances  to  conduct  the  electric  current.  They  convert 
oxygen  into  ozone  and  white  phosphorus  to  the  red  modification. 
Paper  becomes  brown  and  brittle,  the  green  parts  of  plants  lose 
their  coloring  matter,  and  seeds  lose  their  power  to  germinate  under 
the  influence  of  radium  rays.  When  the  skin  is  subjected  to  them 
for  a  short  time  there  is  no  apparent  effect,  but  after  a  week  or  more 
the  portion  which  had  been  exposed  becomes  red  and  inflamed, 
and  behaves  as  though  severely  burned.  The  rays  have  been  used 
with  some  success  for  the  treatment  of  cancer  and  certain  skin 
diseases. 

Water  which  contains  a  dissolved  or  suspended  radium  com- 
pound slowly  separates  into  oxygen  and  hydrogen,  just  as  though 
an  electric  current  was  passing  through  it,  except  that  the  two 
gases  escape  mixed  instead  of  at  different  poles. 


CHAPTER   XXIII 

• 

THE  ZINC  GROUP 
Magnesium,  Zinc,  Cadmium,  Mercury. 

MAGNESIUM,  Mg  =  24. 

240.  Occurrence. — Though  magnesium  itself  does  not 
occur  in  nature,  compounds  of  it  are  common,  and  in  this 
form  it  makes  up  more  than  2J  per  cent  of  the  accessible 
part  of  the  earth's  crust,  standing  sixth  among  the  ele- 
ments   as    regards    abundance.      Magnesium    carbonate, 
MgC03,  is  an  important  mineral,  and  its  double  compound 
with  calcium  carbonate — dolomite,  CaC03MgC03 — is  very 
abundant.    Asbestos  is  a  magnesium  compound,  and  com- 
pounds of  magnesium  are  contained  in  most  common  rocks 
and  soils.     It  is  also  found  in  plants,  especially  in  the 
seeds,   and  in   substances   of   animal   origin,   particularly 
blood,  milk,  and  bones. 

241.  Preparation  and  properties. — The  first  magnesium 
compound  to  be  obtained  was  crystallized  magnesium  sul- 
phate, MgS047H20,  which,  on  account  of  the  mineral  spring 
from  whose  water  it  was  extracted,  is  still  called  "  Epsom 
salts."     The  metal  is   best  prepared   from  the   chloride, 
MgCl2,  by  heating  it  with  sodium : 

MgCl2  +  2Na  =  Mg  +  2NaCl, 

or  by  fusing  it  in  a  crucible  and  passing  the  electric  current 
through  it,  when  chlorine  is  given  off  and  magnesium  left, 
which  is  purified  by  distillation. 
172 


THE  ZINC  GROUP  173 

It  is  a  white  metal,  fairly  malleable  and  ductile,  and 
is  scarcely  acted  upon  by  air  at  ordinary  temperatures. 
Heated  in  the  air  or  in  oxygen,  it  burns  with  a  dazzling 
bluish-white  flame,  producing  a  white  smoke  of  magnesium 
oxide,  MgO.  It  is  doubtful  whether  magnesium  is  able  to 
decompose  water,  but  if  it  does  the  action  is  very  slow. 
However,  we  can  easily  show  that  at  high  temperatures 
its  action  upon  steam  is  violent.  Some  water  is  placed  in 
a  large  beaker  covered  with  an  asbestos  plate  having  a  hole 
in  the  center,  and  the  water  boiled  vigorously  until  the 
air  has  all  been  expelled  from  the  vessel.  Then  a  piece  of 
burning  magnesium  wire — held  in  forceps — is  introduced 
through  the  orifice  into  the  atmosphere  of  steam,  where  it 
goes  on  burning  more  brilliantly  than  in  air.  At  the  same 
time  a  pale  flame  makes  its  appearance  around  the  orifice 
in  the  plate.  This  is  burning  hydrogen  set  free  by  the 
combustion  of  the  magnesium  in  the  interior : 

Mg  +  H20  =  MgO  +  H2. 

In  the  same  way  magnesium  is  able  to  burn  in.  most 
gases  and  vapors  which  contain  oxygen,  combining  with  the 
latter  and  liberating  the  other  constituent.  Thus  it  burns 
in  carbon  dioxide,  carbon  separating  as  soot.  It  removes  the 
oxygen  from  most  metallic  oxides  when  mixed  with  them 
and  heated,  and  some  metals  can  be  obtained  in  this  way 
which  are  not  easily  prepared  by  other  methods. 

242.  TJses. — Magnesium  as  compact  metal  has  not  found 
much  application.  An  alloy  of  10  to  15  per  cent  of  it  with 
aluminium  is  called  magnaUum,  and  has  valuable  properties. 
It  has  a  silver-white,  permanent  luster,  is  very  light,  and 
is  easily  worked  and  cast.  It  is  used  for  scales  on  optical 
instruments  and  for  the  beams  of  balances.  The  light  of 
burning  magnesium  acts  rapidly  upon  the  photographic 
plate,  and  for  this  reason  the  powdered  metal  is  largely 
used  in  the  production  of  flash-light  powders.  For  this 


174  ELEMENTARY  CHEMISTRY 

purpose  it  is  mixed  intimately  with  some  substance  which 
will  supply  oxygen  to  it  and  cause  it  to  burn  instantaneous- 
ly. The  student  can  investigate  this  action  by  mixing  upon 
a  brick  a  little  powdered  magnesium  with  rather  more  than 
its  own  weight  of  potassium  chlorate.  The  mixture  can  be 
ignited  by  thrusting  a  piece  of  paper  into  it  and  setting  fire 
to  it.  When  the  flame  reaches  the  powder  there  is  a  blind- 
ing flash  and  a  cloud  of  white  smoke.  The  magnesium  has 
been  converted  into  its  oxide  by  the  oxygen  of  the  potassium 
chlorate. 

243.  Magnesium  oxide,  MgO,  is  the  product  of  the  burn- 
ing of  the  metal  in  air  or  oxygen.     It  has  been  obtained 
in  cubical  crystals,  but  is  usually  a  white  powder,  not  visibly 
crystalline.     It  can  not  be  completely  melted  in  the  oxy- 
hydrogen  flame,  but  melts  and  boils  at  the  temperature  of 
the  electric  arc. 

It  is  remarkable  as  being  the  only  oxide  which  undergoes  no 
change  when  heated  to  the  temperature  of  the  arc  in  contact  with 
carbon.  With  all  other  oxides  the  oxygen  combines  with  the  carbon 
to  carbon  monoxide,  CO,  and  the  metal  is  liberated,  or  else  combines 
with  more  carbon,  producing  a  carbide.  Thus,  when  zinc  oxide  is 
treated  in  this  way  the  products  are  carbon  monoxide  and  zinc  vapor, 
•while  on  the  other  hand,  when  lime  is  mixed  with  carbon  and 
heated  to  this  temperature,  carbon  monoxide  and  calcium  carbide, 
CaCa,  are  produced.  But  magnesium  oxide  is  not  affected,  except 
that  it  will  melt  and  vaporize  if  the  heating  is  prolonged. 

Magnesium  oxide  is  scarcely  soluble  in  water,  but  dis- 
solves readily  in  acids,  forming  magnesium  salts.  Under 
the  name  "  calcined  magnesia "  it  is  employed  in  medi- 
cine. 

244.  Magnesium  sulphate,  MgS04,  like  many  other  me- 
tallic salts,  can  form  crystals  with  different  amounts  of 
water,  depending  upon  the  temperature  at  which  its  solu- 
tion in  water  is  made  to  crystallize.     The  following  sub- 
stances are  known  in  this  case : 


THE  ZINC  GROUP  175 

MgS04,H20. 

MgS04,6H20. 

MgS04,7H20. 

MgS04,12H20. 

MgS04,24H20. 

Taken  together,  these  are  called  the  hydrates  of  magnesium 
sulphate.  The  forms  of  the  crystals  of  different  hydrates  of 
the  same  salt  are  different,  and  the  rule  holds  good  that 
the  lower  the  temperature  at  which  the  solution  is  made 
to  crystallize  the  more  water  in  the  hydrate  produced. 
Thus,  in  this  case,  the  hydrates  with  12H20  and  24H,0 
can  only  exist  below  0°;  above  that  temperature  they  de- 
compose, even  under  water,  thus : 

MgS04.12H20  =  MgS04.7H20  +  5H20. 

MgS04.7H20  is  obtained  when  the  liquid  is  allowed  to 
crystallize  at  ordinary  temperatures.  When  it  is  heated 
slightly  above  the  boiling-point  of  water  it  decomposes, 
yielding  MgS04.H20— 

MgS04.7H20  =  MgS04.H20  +  6H20. 

And  this,  when  heated  still  higher,  loses  the  last  molec- 
ular weight  of  water,  leaving  the  anhydrous  sulphate 
MgS04. 

It  is  likely  that  all  salts  containing  water  of  crystalliza- 
tion behave  in  a  similar  way.  We  shall  only  have  time  to 
study  the  most  important  hydrate  in  each  case,  which  is 
usually  that  one  formed  when  the  solution  crystallizes  at 
room  temperature,  but  the  student  should  be  careful  not 
to  get  the  false  idea  that  this  is  the  only  hydrate  that  can 
exist. 

Epsom  salts,  MgS04.7H20,  is  white  and  freely  soluble 
in  water.  The  solution  has  a  bitter  taste  and  purgative 
action.  It  is  used  in  medicine  for  this  reason,  and  also  as 
an  antidote  for  lead  poisoning.  It  converts  the  soluble 


176  ELEMENTARY  CHEMISTRY 

lead  compounds  in  the  body  into  lead  sulphate,  which  is  in- 
soluble and  harmless. 

245.  Magnesium  chloride,  MgCl2,  is  contained  in  sea- 
water,  and  in  large  quantities  in  the  water  of  certain  lakes 
— the  Dead  Sea,  for  instance.  In  such  water  marine  life 
is  impossible.  When  magnesium  chloride  is  heated  with 
steam  it  reacts  with  it,  thus : 

MgCl2  +  H20  =  MgO  +  2HC1. 

For  this  reason  care  must  be  taken  not  to  use  water  con- 
taining magnesium  chloride  for  steam  boilers.  The  hydro- 
chloric acid  liberated  in  the  boiler  rapidly  corrodes  it. 

Magnesium  chloride  is  very  soluble  in  water,  and  when 
the  solution  is  evaporated  to  small  bulk,  MgCl2.6H20  sepa- 
rates in  deliquescent  crystals. 


CHAPTER    XXIV 

THE  ZINC  GROUP  (Continued) 

ZINC  AND  CADMIUM. 

ZINC,  Zn  =  65.5. 

246.  Occurrence  and  preparation. — Metallic  zinc  occurs 
in  nature,  but  is  rare.    Two  important  ores  are  the  carbon- 
ate, ZnC03,  and  the  sulphide,  ZnS,  which  is  called  zinc- 
blende.    Traces  of  zinc  compounds  are  found  in  the  human 
liver,  in  eggs,  and  in  the  ashes  of  plants  which  grow  from 
soils  in  which  zinc  is  present. 

Either  zinc  carbonate  or  zinc  sulphide  can  be  converted 
into  zinc  oxide  by  proper  heating.  With  the  carbonate  the 
equation  is — 

ZnC03  =  ZnO  +  C02. 

The  sulphide  must  be  carefully  heated  in  a  current  of 
air,  for  oxygen  takes  part  in  the  process : 

ZnS  +  30  =  ZnO  +  S02. 

The  zinc  oxide  is  then  mixed  with  pounded  coal  and  heated 
to  whiteness: 

ZnO  +  C  =  CO  +  Zn. 

The  zinc  vaporizes  and  is  condensed.  Commercial  zinc  is 
not  pure.  It  contains  traces  of  iron,  lead,  arsenic,  carbon, 
and  other  elements.  Pure  zinc  is  best  obtained  from  it 
by  repeatedly  distilling  it  in  a  vacuum. 

247.  Properties. — Zinc  is  a  bluish-white  lustrous  metal, 
brittle  at  ordinary  temperatures,  but  malleable  when  heated 

177 


178  ELEMENTARY  CHEMISTRY 

to  100°.  At  200°  it  again  becomes  brittle.  It  is  about 
seven  times  as  dense  as  water,  and  melts  below  a  red 
heat.  When  zinc  is  heated  above  its  meltifig-pomt  in  a 
covered  crucible  it  takes  fire,  when  the  cover  is  removed, 
and  burns  with  a  blue-white  flame,  filling  the  air  with 
white  flakes  of  zinc  oxide.  We  have  seen  (p.  36)  that  com- 
mercial zinc  readily  dissolves  in  dilute  hydrochloric  or 
sulphuric  acid,  and  that  this  is  a  convenient  method  of 
making  hydrogen.  Pure  zinc,  on  the  contrary,  is  not 
affected  by  either  acid,  but  if  a  little  finely  divided  platinum 
or  copper  is  present  with  the  zinc,  the  latter  is  rapidly 
attacked. 

Zinc  is  largely  employed  in  the  manufacture  of  brass 
and  for  the  extraction  of  silver  from  lead  (p.  152).  Gal- 
vanized iron  is  simply  iron  plated  with  zinc.  The  surface 
of  the  iron  is  carefully  cleansed,  and  it  is  then  dipped  into 
molten  zinc  and  allowed  to  cool. 

248.  Zinc  oxide,  ZnO,  can  be  made  by  burning  the  metal 
in  the  air.    In  making  it  on  the  large  scale,  low-grade  zinc 
ores  are  heated  with  coal,  and  the  zinc  vapor,  instead  of 
being  condensed,  is  allowed  to  burn  and  collected  as  zinc 
oxide.    It  is  a  white  powder,  which  becomes  sulphur  yellow 
when  heated,  and  returns  to  its  original  white  when  cold. 
It  is  insoluble  in  water,  but  dissolves  readily  in  acids.    Un- 
der the  name  "  zinc  white/'  it  is  largely  used  as  a  basis  of 
white  paint. 

249.  Zinc  sulphide,  ZnS. — When  powdered  zinc  is  mixed 
with  sulphur  and  a  flame  applied,  the  mixture  burns  explo- 
sively, producing  a  dense  white  smoke   of  zinc   sulphide. 
Zinc  sulphide  occurs  in  nature  as  zinc-blende,  which  is  some- 
times colorless  and  transparent,  but  usually  black  or  brown, 
owing  to  the  presence  of  iron  sulphide.    It  is  the  most  im- 
portant ore  of  zinc.     When  soluble  sulphides  are  added  to 

++ 
solutions  of  zinc  salts — that  is,  when  S  ions  and  Zn  ions  are 


THE  ZINC  GROUP  179 

brought  together  —  zinc  sulphide  is  obtained  as  a  pure  white 
precipitate,  e.  g.  : 

ZnCl2  +  Na2S  =  ZnS  +  2NaCl. 

250.  Zinc  sulphate,  ZnS04,  can  be  obtained  by  dissolv- 
ing zinc  or  zinc  oxide  in  sulphuric  acid.  Several  hydrates 
exist,  of  which  ZnS04.7H20  is  best  known.  It  is  used  in 
medicine  as  a  violent  emetic.  Wood  soaked  in  zinc  sulphate 
solution  is  protected  to  some  extent  both  from  fire  and 
decay. 

The  compounds  of  zinc  with  the  halogens  *  are  all  color- 
less and  very  soluble  in  water.  Zinc  chloride,  ZnCl2,  which 
can  be  obtained  by  burning  the  metal  in  chlorine,  or,  better, 
by  dissolving  it  in  hydrochloric  acid,  is  the  most  important. 
Its  solution  in  water  is  used  as  an  embalming  fluid,  and  also 
for  cleansing  the  surface  of  metals  before  soldering. 

Soluble  zinc  compounds  are  poisonous,  the  chloride  being 
the  most  dangerous.  Since  water,  in  presence  of  air,  has  a 
decided  action  upon  zinc,  it  ought  not  to  be  passed  through 
pipes  or  preserved  in  vessels  lined  with  the  metal.  Work- 
men about  zinc  furnaces  sometimes  suffer  from  a  kind  of 
slow  poisoning,  but  this  may  be  partly  due  to  the  arsenic 
which  ores  of  zinc  usually  contain. 


CADMIUM,  Cd  = 

251.  Cadmium  occurs  in  nature  along  with  zinc,  which 
it  strongly  resembles.  It  is  a  tough,  white  metal,  not  very 
hard,  and  slightly  denser  than  zinc. 

Many  alloys  of  cadmium  have  very  low  melting-points, 
and  it  is  employed  in  the  production  of  fusible  metal,  which 
is  an  alloy  of  tin,  lead,  bismuth,  and  cadmium,  and  melts 
easily  below  the  boiling-point  of  water.  This  alloy  is  used 

1  The  four  elements  fluorine,  chlorine,  bromine,  and  iodine  resem- 
ble each  other  greatly.     Taken  together,  they  are  called  the  halogens. 
13 


180  ELEMENTARY  CHEMISTRY 

in  the  construction  of  automatic  fire-extinguishing  appara- 
tus. Tubes  supplying  water  are  closed  with  plugs  of  fusible 
metal,  so  that  the  heat  of  the  fire  melts  the  plug  and  re- 
leases the  water. 

Cadmium  oxide,  CdO,  is  the  product  of  the  burning  of 
the  metal  in  the  air  or  oxygen.  It  is  a  brown,  crystalline 
powder. 

Cadmium  sulphide,  CdS,  occurs  in  nature,  and  can  be 
made  by  allowing  hydrogen  sulphide  gas  to  bubble  through  a 
water  solution  of  cadmium  chloride: 

CdCl2  +  H2S  =  CdS  +  2HC1. 

Under  the  name  "  cadmium  yellow  "  it  is  much  used  as  a 
color  in  oil  paints  and  for  coloring  toilet  soaps.  It  is  a 
bright  yellow  powder,  insoluble  in  water. 


CHAPTER    XXV 

MERCURY 
Hg  =  200. 

252.  Historical. — Mercury  was  known  to  the  ancients. 
All  through  the  middle  ages,  metals  which  we  now  know 
to  be  elements — copper,  lead,  silver,  and  gold,  for  example — 
were  regarded  as  compounds  of  mercury  with  other  constitu- 
ents, chiefly  sulphur  and  salt.     It  was  supposed  to  be  the 
mercury  in  them  which  gave  them  their  luster.    This  base- 
less notion  was  responsible  for  the  efforts  of  the  alchemists 
— lasting  more  than  a  thousand  years  and  not  yet  entirely 
over — to  prepare  silver  and  gold  from  the  other  metals. 

253.  Occurrence  and  preparation. — Mercury  occurs  na- 
tive in  drops  distributed  through  slates   and  sandstones, 
but  the  chief  ore  is  cinnabar,  mercuric  sulphide,  HgS.    This 
occurs  in  dark-red,  lustrous  crystals,  and  is  abundant  in  a 
few  localities — e.  g.,  in  Spain,  California,  and  Mexico.    For 
the  extraction  of  the  metal,  the  sulphide  is  roasted  in  a  cur- 
rent of  air.    When  most  sulphides  are  treated  in  this  way, 
sulphur  dioxide,  S02,  and  the  oxide  of  the  metal  are  pro- 
duced.    But  mercuric  oxide  is  decomposed  by  heat,  and 
therefore  the  metal  is  formed : 

HgS  +  02  =  S02  +  Hg. 

The  mercury  condenses  in  stone  chambers,  through  which 
the  gases  are  led. 

254.  Properties. — Mercury  is  one  of  the  two  liquid  ele- 
ments, bromine  being  the  other.    It  has  a  strong  silver-white 

181 


182  ELEMENTARY  CHEMISTRY 

metallic  luster  which  is  permanent  in  the  air.  At  — 39° 
it  freezes  to  a  lustrous  metallic  mass  which  can  be  beaten 
out  under  the  hammer  like  lead.  Liquid  mercury  vaporizes 
slowly  at  ordinary  temperatures,  and  if  some  be  placed  in 
the  bottom  of  a  bottle  and  a  piece  of  gold  leaf  suspended 
from  the  cork  above  it,  the  gold  will  slowly  turn  white  from 
mercury  deposited  upon  it.  At  357°,  under  one  atmosphere 
pressure,  mercury  boils,  passing  into  a  colorless  vapor.  The 
density  of  this  vapor,  referred  to  hydrogen,  is  100.  There- 
fore the  molecular  weight  must  be  100X2  =  200.  Since 
the  atomic  weight  is  also  about  200,  the  molecule  of  mer- 
cury can  contain  but  one  atom ;  the  atom  and  the  molecule 
are  identical.  The  same  thing  is  true  of  other  metals  in  the 
state  of  vapor,  and,  we  have  reason  to  think,  in  the  solid 
and  liquid  state  also. 

Mercury  is  not  affected  by  water  at  any  temperature,  nor 
by  oxygen  or  air  in  the  cold.  Heated  just  below  its  boil- 
ing-point in  the  air  it  slowly  passes  into  a  red  powder 
of  mercuric  oxide,  HgO.  The  metal  is  easily  dissolved  by 
nitric  acid,  which,  when  cold  and  dilute,  produces  mer- 
curous nitrate,  Hg2(N03)2;  when  hot  and  stronger,  mer- 
curic nitrate,  Hg(N03)2. 

255.  From  the  last  sentence  it  will  be  seen  that  there  are  two 
sets  of  salts  of  mercury.  For  example : 

Mercurous  salts.  Mercuric  salts. 

Chloride HgaCla  HgCl2 

Nitrate Hg2(NO3)2  Hg(NO,)2 

Sulphate HgaSO4  HgSO4 

The  structure  of  the  mercuric  salts  is  simple,  the  mercury  being 
bivalent,  e.  g. : 

Cl 
Mercuric  chloride,  Hg<^| 

In  the  mercurous  compounds  we  think  of  the  two  mercury  atoms  as 
linked  together,  e.  g. : 

Hg-Cl 
Mercurous  chloride,    I 

Hg-CL 


MERCURY  183 

The  mercury,  then,  is  still  bivalent  in  the  mercurous  compounds. 
In  solutions  of  mercuric  compounds  the  mercury  exists  as  ions  Hg, 

Hg 
while  in  mercurous  solutions  the  ion  is    |   .     The  chemical  behavior 

+ 

of  these  two  ions  is  very  different. 

The  student  should  never  accept  doubled  formulas  on  faith,  but 
should  form  the  habit  of  inquiring  what  facts  cause  us  to  double 
them.  When  we  do  not  know  what  the  molecular  weight  of  a  com- 
pound is,  it  is  proper  to  use  the  simplest  formula  which  will  express 
its  chemical  composition.  This  we  do,  for  instance,  with  substances 
like  silver  chloride,  AgCl,  and  mercuric  oxide,  HgO,  whose  moleculr.r 
weights  are  unknown  at  present.  But  when  we  double  the  formula 
we  must  have  facts  which  support  us  in  so  doing  if  the  formula  is 
to  have  any  meaning.  In  this  case  the  facts  are  these :  the  density 
of  the  vapor  into  which  pure  dry  mercurous  chloride  is  converted  by 
heat  is  about  235.5;  therefore,  the  molecular  weight  is  235.5  x  2,  or 
471,  and  the  formula  must  be  Hg2Cla.  Further,  the  molecular  weight 
of  mercurous  nitrate  dissolved  in  water  (p.  124)  is  found  to  agree 
with  the  formula  Hga(NO3)2,  not  HgNO3. 

256.  Mercurous  compounds. — Mercurous  chloride,  Hg2Cl2, 
is  largely  used  in  medicine  under  the  name  "  calomel."  It 
can  be  made  by  heating  gently  an  intimate  mixture  of  mer- 
curic chloride  and  mercury  when  it  vaporizes  and  is  con- 
densed : 

HgCl2  +  Hg  =  Hg2Cl2. 

It  is  white,  crystalline,  and  insoluble  in  water.  When 
heated,  it  vaporizes  without  melting.  Like  silver  chloride, 
it  darkens  on  exposure  to  light. 

Mercurous  nitrate,  Hg2(N03)2,  is  the  most  important 
soluble  mercurous  salt.  It  is  produced  when  mercury  dis- 
solves in  cold  dilute  nitric  acid. 

257.  Mercuric  compounds. — Mercuric  oxide,  HgO,  is  ob- 
tained by  heating  mercuric  nitrate: 

Hg(N03)2  =  HgO  +  2N02  +  0. 


184  ELEMENTARY  CHEMISTRY 

It  is  a  dense,  red,  crystalline  powder,  which  slowly  decom- 
poses in  light  into  mercury  and  oxygen.  The  color  of  mer- 
'curic  oxide,  like  that  of  many  other  substances,  depends  upon 
the  temperature.  At  —  200°  (the  temperature  of  boiling 
liquid  air)  it  is  sulphur-yellow,  at  room-temperature  red, 
at  higher  temperatures  black.  Below  a  red  heat  it  separates 
into  mercury  and  oxygen  (p.  26). 

Mercuric  sulphide,  HgS,  is  the  mineral  cinnabar.  A  bril- 
liant scarlet  variety  of  it,  called  vermilion,  is  made  by 
grinding  mercury  and  sulphur  together  in  presence  of  a  so- 
lution of  potassium  sulphide,  K2S.  This  is  much  used  as  a 
coloring  matter.  The  dense  black  precipitate  which  is 
formed  when  hydrogen  sulphide  is  passed  into  a  solution  of 
a  mercuric  salt  is  mercuric  sulphide. 

Mercuric  sulphide  approaches  absolute  insolubility  in 
water  more  nearly  than  any  other  compound  which  has  been 
investigated  thus  far.  The  cinnabar  of  nature  ranges  in 
color  from  nearly  black  to  bright  red  ;  the  shade  of  vermilion 
is  much  affected  by  the  method  of  preparation,  and  the  pre- 
cipitated HgS  is  invariably  black.  These  differences  in 
color  are  remarkable,  and  not  yet  satisfactorily  explained. 

Mercuric  chloride,  HgCl2,  is  produced  when  the  metal 
burns  in  chlorine.  It  can  be  made  by  grinding  mercuric 
sulphate  with  salt  and  then  heating  the  mixture  in  large 
glass  flasks.  The  vapor  of  mercuric  chloride  condenses  to 
a  white  compact  mass  in  the  upper  cool  portion  of  the  flask, 
which  is  broken  in  order  to  remove  it  : 


HgS04  +  SNaCl  =  HgCl2  +  Na2S04. 

Mercuric  chloride  is  a  dense,  white,  crystalline  salt,  which 
melts  first,  and  then  vaporizes,  when  heated.  It  is  freely 
soluble  in  water,  and  the  solution  is  almost  a  non-con- 
ductor of  the  current,  showing  that,  unlike  most  other  salts, 
mercuric  chloride  scarcely  separates  into  its  ions  when  dis- 
solved. It  is  intensely  poisonous,  and  the  antidote  is  some 


MERCURY  185 

form  of  albumin — for  instance,  raw  eggs  taken  at  once — 
for  albumin  forms  with  it  an  insoluble  compound.  Mer- 
curic chloride  is  one  of  the  best  of  all  antiseptics  and  dis- 
infectants, and  the  dilute  water  solution  is  largely  used  for 
this  purpose  in  surgery. 

Mercuric  iodide,  HgI2,  separates  when  a  potassium  iodide 
solution  is  added  to  one  of  mercuric  chloride  as  a  precipitate 
which  is  at  first  yellow,  but  turns  red  at  once.  An  excess 
of  either  potassium  iodide  or  mercuric  chloride  will  redis- 
solve  it.  Mercuric  iodide  is  an  intensely  red  crystalline  pow- 
der. We  have  already  discussed  its  behavior  when  heated 
(p.  24).  It  is  insoluble  in  water,  but  soluble  in  alcohol  and 
in  water  containing  potassium  iodide. 

258.  The  amalgams. — This  name  is  given  to  the  alloys 
which  mercury  forms  with  other  metals.  They  are  very  in- 
teresting because  they  are  probably  similar  in  character  to 
other  alloys,  and,  owing  to  the  fact  that  mercury  is  a  liquid, 
they  are  much  easier  to  make  and  to  work  with.  Hence  we 
may  expect  their  study  to  throw  light  upon  the  nature  of 
alloys  in  general — an  interesting  question.  The  amalgams 
are  silver-white,  crystalline,  metallic  solids.  They  are  some- 
times said  to  be  pasty  masses,  but  this  is  because  too  much 
mercury  was  used  in  making  them  and  the  soft  product  was 
a  mixture  of  the  real  amalgam  with  liquid  mercury. 

Mercury  usually  forms  several  amalgams  with  the  same  element. 
Thus,  with  sodium,  the  compounds  NaHg5  and  NaHge  are  known, 
and  with  potassium,  the  compounds  KHgio,  KHg13,  and  KHgu. 
KHgi4  can  only  exist  below  0°,  at  which  temperature  it  separates 
into  KHgia  and  mercury — 

KHgu  =  KHg12  +  2Hg. 

From  0°  to  73°  KHgu  exists.  At  the  latter  temperature  it  sepa- 
rates, thus: 

KHg13  =  KHg10  +  2Hg. 

So  with  other  amalgams.  Those  richer  in  mercury  can  only  exist  at 
low  temperatures,  and  when  heated,  split  up,  yielding  mercury  ?»nd 
amalgams  poorer  in  it.  It  will  be  seen  that  the  mercury  in  the 


186  ELEMENTARY  CHEMISTRY 

amalgams  behaves  very  much  in  the  same  way  as  the  water  in  salts 
containing  water  of  crystallization  (p.  175). 

259.  General  remarks  on  the  calcium  and  zinc  groups. — 

The  relation  between  the  calcium  group  and  the  zinc  group 
is  similar  to  that  between  the  sodium  group  and  the  copper 
group.  Like  the  sodium  group,  the  calcium  group  is  com- 
posed of  strong  metals — using  the  word  in  its  chemical  sense 
— which  are  energetically  acted  upon  by  water,  forming  their 
hydroxides,  and  these  hydroxides  are  strong  bases.  Their 
susceptibility  to  the  action  of  water  prevents  calcium,  stron- 
tium, and  barium  from  receiving  any  application  as  metals. 

On  the  other  hand,  the  metals  of  the  zinc  group  are 
either  slightly  or  not  at  all  affected  by  water  and  air,  and 
are  accordingly  employed  in  the  metallic  state  for  various 
purposes.  Mercury,  like  gold,  is  not  affected  by  the  atmos- 
phere at  all,  and  will  retain  its  luster  unchanged  for  any 
length  of  time. 

What  is  the  chemical  difference  between  the  "  noble  " 
metals,  like  silver,  gold,  and  platinum,  and  the  so-called 
"  base  "  metals,  like  lead,  iron,  and  zinc  ?  The  difference  is 
simply  one  of  chemical  activity.  Gold  is  an  inert  element. 
Its  tendency  +o  enter  into  chemical  combination  with  other 
elements  is  small,  and  hence  its  luster  is  unaffected  by  the 
atmosphere  even  at  high  temperatures.  Iron,  on  the  con- 
trary, possesses  a  much  stronger  tendency  to  take  part  in 
chemical  reactions.  Therefore,  at  the  surface  of  a  piece  of 
the  metal  there  begins  a  chemical  change  in  which  the  water 
and  oxygen  of  the  air  convert  the  surface  of  the  iron  into 
a  reddish-brown  hydroxide  called  "  rust."  And,  in  general, 
the  chief  distinction  between  the  "  noble  "  metals  and  the 
common  metals  is  that  the  latter  are  far  more  active  chem- 
ically. We  must  remember  also  that  the  noble  metals  are 
not  very  abundant  in  nature,  and  are  accordingly  high  in 
price.  The  metals  of  both  the  calcium  and  the  zinc  groups 
are  almost  always  bivalent. 


CHAPTER    XXVI 

BORON  AND  ALUMINIUM 
Boron,  B.  Aluminium,  Al. 

260.  Boron  is  unmistakably  a  non-metal.     It  has  none 
of  the  physical  properties  of  a  metal,  it  forms  a  gaseous 
hydrogen  compound  (BH3),  yields  no  salts  with  acids,  and 
its  hydroxide  is  an  acid.    Aluminium  is  physically  a  metal, 
and  for  the  most  part  chemically  also,  though  its  hydroxide 
is  only  a  feeble  base.    Both  elements  are  almost  always  tri- 
valent. 

BORON,  B  =  11. 

261.  Occurrence. — Boron  occurs  in  nature  as  boric  acid 
or  salts  of  boric  acid.     Boric  acid  itself,  B(OH)3  (4»ornn 

jiydt^TtJS^,  is  contained  in  the  steam  which  issues  from  the 
ground  in  volcanic  regions,  particularly  in  Tuscany. 
Around  such  a  steam-jet  a  stone  basin  is  built  in  such  a 
way  that  the  steam  is  made  to  bubble  through  water  in  the 
basin.  The  bor(te^hyclrox<MO  dissolves  in  this  water,  which, 
after  being  treated  with  steam  in  several  such  basins,  is 
evaporated  to  recover  it.  For  this  purpose  the  heat  of 
the  volcanic  steam  is  employed,  so  as  to  save  the  expense 
of  artificial  fuel.  The  liquid  is  made  to  flow  slowly  over 
a  long  inclined  lead  plate,  which  is  heated  by  steam 
passing  beneath  it.  Borax,  N~a2B407.10H20,  is  largely  ob- 
tained from  the  water  of  certain  lakes  in  California,  Cey- 
lon, and  elsewhere. 

187 


188  ELEMENTARY  CHEMISTRY 

262.  Preparation  and  properties.  —  Boron  can  be  ob- 
tained by  heating  boron  oxide  B203,  with  sodium  or  mag- 
nesium : 

B203  +  6Na  =  3Na20  +  2B, 

or  by  heating  potassium  in  boron  fluoride,  which'  is  a  gas: 
BF3  +  3K  =  B  +  3KF. 

It  is  a  brown  powder,  odorless  and  tasteless,  and  slightly 
soluble  in  water.  Heatedln  the  air  or  in  oxygen,  it  burns 
brightly  to  B203.  At  a  high  temperature  it  combines  with 
nitrogen  when  heated  in  the  gas,  producing  boron  nitride, 
BN",  a  white  powder. 

Boron  oxide,  B203,  is  made  by  strongly  heating  the  hy- 
droxide — 


It  is  a  colorless,  brittle,  glassy  mass,  which  melts  readily, 
but  only  vaporizes  at  very  .high  temperatures.  * 

The  preparation  of  borm  /w^&^^fe/  B(OH)3,  has  just 
been  described.  It  forms  colorless  crystals  which  are  solu- 
ble in  water  and  in  alcohol.  The  latter  solution  burns 
with  a  beautiful  green  flame,  and  this  peculiarity  is  used 
as  a  test  for  boron  hydroxide.  Borb^fe^oite^ide  vis  feebly 
acid  to  litmus,  and  reacts  with  bases  forming  salts  in  which 
its  hydrogen  is  replaced  by  various  metals.  It  is  therefore 
called  boric  acid,  and  its  salts  are  called  the  borates.  The 
dilute  water  solution  of  boric  acid  forms  an  excellent  wash 
for  sore  or  inflamed  eyes.  Large  quantities  of  the  acid 
are  used  for  the  manufacture  of  borax. 

263.  Borax,  Na2B407.10H20,  forms  colorless  crystals, 
which  give  off  flashes  of  light  when  crushed  in  the  dark. 
It  is  freely  soluble  in  water,  and  the  solution  has  a  peculiar 
cooling  taste. 

When  borax  is  strongly  heated  the  water  of  crystalliza- 
tion escapes  and  the  liquid  which  is  loft  solidifies  on  cool- 
ing to  a  hard,  brittle,  glassy  mass  called  borax  glass,  which 


FEIEDRICH   WOHLER 
B.  Germany,  1800.     D.  1882. 


BORON  AND  ALUMINIUM  189 

is  simply  Na2B407.  Borax  glass  dissolves  many  compounds 
of  the  metals,  and  is  often  colored  by  them  in  such  a  way 
that  the  metal  can  be  identified.  This  behavior  is  used 
in  analysis.  Thus,  if  we  heat  a  little  borax  on  a  platinum 
wire,  it  swells  up,  and  finally,  when  the  steam  has  all  es- 
caped, solidifies  on  cooling  to  a  colorless  bead  of  borax 
glass.  If,  now,  we  cause  a  fragment  of  cobalt  nitrate  (or 
any  other  cobalt  compound)  to  adhere  to  the  bead,  and  heat 
it  again,  it  becomes  deep  sapphire  blue.  If,  instead  of  a 
cobalt  salt,  we  use  a  compound  of  manganese,  the  bead  is 
colored  amethyst. 

When  carbon  and  boron  are  heated  together  to  a  very 
high  temperature,  they  combine,  producing  CB6,  a  black 
crystalline  compound,  which  is  interesting  on  account  of 
its  extreme  hardness.  It  is  about  as  hard  as  the  diamond, 
and  its  powder  can  be  used,  instead  of  diamond  dust,  for 
cutting  the  gem. 

ALUMINIUM,  Al  =  27. 

264.  Occurrence. — Aluminium   does  not   occur  native, 
but  its  compounds  are  plentiful.     In  abundance  it  is  third 
among  the  elements,  making  up  nearly  8  per  cent  of  the 
earth's  crust.    Particularly  common  are  compounds  of  alu- 
minium with  silicon  and  oxygen,  called  silicates   of  alu- 
minium.   Clay  is  a  compound  of  this  class,  having,  in  pure 
condition,  the  composition  Al2Si207  with  water.    The  vari- 
ous felspars — most  important  rock-forming  minerals — are 
also  aluminium  silicates,  but  they  contain  other  metals,  par- 
ticularly sodium,  potassium,  and  calcium.     Cryolite,  which 
occurs  abundantly  in  Greenland,  is  a  fluoride  of  aluminium 
and  sodium,  AlF33NaF.    It  is  employed  in  the  manufacture 
of  aluminium,  and  also,  on  account  of  its  sodium,  in  the 
production  of  sodium  carbonate. 

265.  Preparation. — Aluminium  was  first   obtained  by 
heating  the  chloride  with  potassium,  and  for  a  long  time  a 


1(JO  ELEMENTARY  CHEMISTRY 

similar  method  was  used  in  preparing  it.  It  is  now  made 
on  the  large  scale  by  an  electrolytic  process  which  has  great- 
ly reduced  its  price. 

The  mineral  cryolite  melts  at  a  low  temperature  and 
the  liquid  dissolves  aluminium  oxide,  A1203.  When  the  cur- 
rent from  a  dynamo  is  passed  through  the  solution,  only 
the  A1203  is  decomposed,  the  cryolite  not  being  affected. 
In  practice  cryolite  is  melted  in  a  vessel  lined  with  carbon, 
and  aluminium  oxide  added.  Then  a  bundle  of  carbon 

rods  is  connected  with  the 
positive  pole  of  the  cur- 
rent from  a  dynamo  and 
dipped  into  the  liquid. 
The  vessel  itself  is  con- 
nected with  the  negative 
pole.  When  the  current 
passes,  the  aluminium, 
like  all  metals,  goes  to 
the  negative  pole  and  col- 
lects, in  the  melted  state, 

FIG.  35.-Industrial  production  of  aluminium    be]ow     th       melted     Crvo- 
by  electrolysis.  .  ~ 

lite.  From  time  to  time 

it  is  withdrawn  and  fresh  aluminium  oxide  added,  so  that 
the  process  is  continuous  (Fig.  35). 

266.  Properties. — Powdered, aluminium  is  gray ;  the  com- 
pact metal  has  about  the  color  aud  luster  of  tin.  It  is  very 
light,  the  density  being  about  that  of  glass — 2.7 — and 
very  malleable  and  ductile.  Since  its  tenacity  is  only  about 
i  that  of  steel,  it  has  no  chance  of  displacing  the  latter 
for  structural  purposes. 

Aluminium  melts  at  a  clear  red  heat  to  a  thin  liquid, 
and  vaporizes  only  at  the  very  highest  temperatures.  It  is 
not  acted  upon^  by  either  dry  or  moist  air,  and  in  compact 
masses  not  much  affected  by  beinglieafetHn  air  or  oxygen. 
As  powder  it  burns  brightly  when  heated.  Aluminium  is 


BORON  AND  ALUMINIUM  191 

not  acted  upon  by  nitric  acid  nor  by  dilute  sulphuric  acid. 
It  dissolves  readily  in  hydrochloric  acid  and  in  solutions  of 
the  hydroxides  of  potassium  or  sodium,  hydrogen  escaping 
in  both  reactions.  The  metal,  especially  in  the  presence  of 
salt,  is  somewhat  readily  attacked  by  the  acids  which  are 
present  in  fruits  and  vegetables,  and  has  not  proved  very  suc- 
cessful as  a  material  for  cooking  utensils. 

267.  Uses. — Flash-light  powders  containing  aluminium 
have  recently  appeared,  and  if  they  are  successful  the  metal 
will  entirely  displace  magnesium  for  this  purpose,  for  it  is 
much  cheaper.    When  a  little  aluminium  is  added  to  melted 
steel  just  before  it  is  poured  into  the  mold,  the  casting 
obtained  is  solid  and  free  from  "  blow-holes  " — little  cav- 
ities which  have  a  very  bad  effect  on  the  strength  of  the 
steel.     The  metal  is  largely  used  for  this  purpose. 

At  a  high  temperature,  aluminium  acts  upon  metallic 
oxides,  combining  with  the  oxygen  and  liberating  the  metal, 
e.  g.: 

Cr203  +  2A1  =  A1203  +  2Cr. 

Much  energy  is  evolved  in  the  change,  and  an  extremely 
high  temperature  (3000°)  produced.  Manganese  and 
chromium  are  now  made  commercially  by  this  method. 

Many  alloys  containing  aluminium  have  been  made. 
The  most  important  is  aluminium  bronze,  which  contains 
5  to  8  per  cent  of  aluminium,  the  rest  being  copper.  It 
has  about  the  color  of  gold,  is  very  rigid  and  strong,  and  is 
unaffected  by  air  and  water.  Its  properties  make  it  valu- 
able for  many  purposes. 

268.  Aluminium  oxide,  A1203,  is  the  mineral  corundum. 
When  perfectly  transparent  and  free  from  flaws  itiorms, 
when  tinted  red  by  the  presence  of  a  little  chromium,  the 
ruby;  when  blue  by  traces  of  cobalt,  the  sapphire.     Both 
of  these  gems  have  been  made  artificially.    An  impure  form 
.of  corundum,  dark  from  the  presence  of  iron  oxide,  is  called 


192  ELEMENTARY  CHEMISTRY 

emery.  Aluminium  oxide  can  be  made  artificially  by  burn- 
ing the  metal  in  the  air  or  by  heating  the  hydroxide.  Crys- 
tallized aluminium  oxide  is  intensely  hard.  Among  min- 
erals, it  stands  next  to  the  diamond  in  this  respect,  but 
some  artificial  products  are  known  which  must  be  placed 
between  the  two. 

269.  Aluminium  hydroxide,  A1(OH)3  is  precipitated  in 

+++    . 
gray  flakes  when  aluminium  ions,  Al,  and  hydroxyl  ions 

come  together — i.  e.,  when  a  solution  of  a  base,  like  NaOH, 
is  added  to  a  solution  of  an  aluminium  salt,  like  A1C18. 
It  is  soluble  in  acids  producing  aluminium  salts,  and  solu- 
ble also  in  solutions  of  bases  producing  compounds  called 
aluminates  in  which  its  own  hydrogen  is  replaced  by 
metals,  e.  g. : 

A1(OH)S  +  3NaOH  =  Al(ONa)3  +  3H20. 

Al(ONa)3  is  sodium  aluminate.  Thus,  A1(OH)3  acts 
like  a  base  with  acids  and  like  an  acid  with  bases.  Heat 
converts  aluminium  hydroxide  into  the  oxide: 

2A1(OH)3  =  A1203  +  3H20. 

270.  Aluminium  sulphate,  A12(S04)3,  can  be  made  by 
boiling  clay  with  sulphuric  acid.    It  forms  pearly  scales  con- 
taining 18H20,  which  are  freely  soluble  in  water.     It  is 
largely  used  in  dyeing,  not  as  a  color,  for  it  is  colorless,  but 
for  the  purpose  of  causing  dyestuffs  to  adhere  to  the  fabric. 
Cotton,  in  the  presence  of  aluminium  sulphate,  is  permanent- 
ly colored  by  many  dyes  which  do  not  affect  it  when  alone. 
A  substance  which  behaves  in  this  way  is  called  a  mordant 
by  the  dyer.    Formerly  mordants  were  universally  employed 
in  dyeing  cotton  fabrics,  but  now  numerous  dyes  are  known 
which  color  cotton  perfectly  without  their  aid. 

Ordinary  ALUM  is  a  double  sulphate  of  potassium  and 
aluminium : 

,  K.SO.. 


BORON  AND  ALUMINIUM 


193 


When  it  has  been  crystallized  from  water  it  contains, 
in  addition,  24H20,  so  that  the  composition  of  alum  in 
crystals  is 

A12(S04)3.K2S04.24H20. 

This  is  potassium  alum.    Many  univalent  metals  form  sim- 
ilar alums.     Thus  sodium  alum  is 

Al2(S04)3.Na2S0424H20. 

Ammonium,  NH4,  which  acts  so  much  like  sodium  in 
other  respects,  also  forms  an  alum: 

A12(S04)3(NH4)2S0424H20. 

This  is  ammonium  alum,  and  is  an  important  alum  of  com- 
merce. 

Further,  we  may  have  other  trivalent  metals  in  place 
of  aluminium  in  an  alum.  Thus,  chromium  alum,  known 
in  trade  as  "  chrome  alum," 
is 

Cr2(S04)3K2S04.24H20, 

and,  just  as  with  the  alumin- 
ium alums,  so  here,  other 
univalent  metals  can  take 
the  place  of  the  potassium. 
The  total  number  of  differ- 
ent alums  is  therefore  very 
great.  They  are  very  simi- 
lar to  each  other,  all  crys- 
tallizing in  the  same  form 

271.  Two  or  more  substances  which  crystallize  in  the  same  form 
are  said  to  be  isomorphous,  and  isomorphism  is  frequently  connect- 
ed, as  it  is  here  among  the  alums,  with  a  close  similarity  in  chemi- 
cal make-up.  This  fact  is  a  valuable  help  at  times  in  determining 
the  atomic  weight  of  an  element.  Thus,  suppose  the  atomic  weight 
of  rubidium  to  be  unknown.  We  make  the  rubidium  alum,  whose 
formula,  if  it  is  similar  to  the  potassium  alum,  must  be 


PIG.  36.— Alum  crystals. 

(Fig.    36). 


194  ELEMENTARY  CHEMISTRY 

Now  the  amount  of  potassium  in  a  molecular  weight  of  the  potas- 
sium alum  is  39  X  2  =  78  parts ;  and  when  we  analyze  the  rubidium 
alum  we  find  that  it  requires  170  parts  of  rubidium  to  play  the  same 
role.  But  since  the  two  compounds  are  isomorphous,  their  formulas 
are  in  all  probability  similar,  and  these  170  parts  are  probably  two 
atomic  weights,  Rba.  Hence  the  atomic  weight  of  rubidium  is 

probably  —  =  85. 

The  alums  can  be  made  by  bringing  solutions  of  the 
two  single  sulphates  together  when  they  crystallize  out,  be- 
ing less  soluble  in  water  than  the  salts  singly.  Thus,  when 
strong  potassium  sulphate  and  strong  aluminium  sulphate 
solutions  are  mixed,  potassium  alum  separates  in  crystals. 

272.  Aluminium  chloride,  A1C13,  is  best  made  by  heating 
cuttings  of  the  metal  to  redness  in  a  wide  glass  tube,  while 
a  rapid  current  of  hydrochloric  acid  gas  is  passed  over 
them.  It  forms  colorless  crystals,  which  rapidly  dissolve 
in  water,  the  liquid  becoming  warm.  The  water  solution 
is  acid  to  litmus  and  other  indicators,  and  therefore  must 
contain  hydrogen  ions.  As  a  matter  of  fact,  when  alumin- 
ium chloride  is  dissolved  in  water,  a  reaction  occurs  which 
liberates  hydrochloric  acid : 

A1C13  +  3H20  =  A1(OH)8  +  3HC1. 

+ 
Since  the  hydrochloric  acid  is  dissociated  into  H  and 

Cl  ions,  the  liquid  is  acid.  Whenever  the  solution  of  a 
salt  is  found  to  be  acid,  a  reaction  of  this  sort  must  have 
occurred  between  the  salt  and  the  water  (p.  127). 


CHAPTER   XXVII 

THE  CARBON  GROUP 
Carbon,  C.         Silicon,  Si.        Tin,  Sn.        Lead,  Pb. 

273.  The  elements  of  this  group  are  usually  quadriva- 
lent, though  tin  and  lead  form  many  compounds  in  which 
they  are  bivalent. 

Carbon  and  silicon  are  non-metals;  tin  and  lead  are 
metals. 

Carbon  is  such  an  important  element  that  we  shall  leave 
it  to  the  last,  and  study  it  somewhat  more  in  detail  than  the 
others. 

SILICON,  Si  =  28.5. 

274.  Occurrence. — Silicon,  although  it  never  occurs  free 
in  nature,  stands  next  to  oxygen  in  abundance,  making  up 
more  than  one-fourth  of  the  rock-masses  of  the  earth's 
crust.    Its  oxide,  Si02 — called  quartz — is  the  most  common 
of  minerals.     Compounds  of  the  metals  with  silicon  and 
oxygen,  called  silicates,  make  up  the  great  bulk  of  most 
important  rocks,  like  granite,  gneiss,  and  serpentine.     As- 
bestos is  a  silicate  of  magnesium,  MgSi03,  usually  con- 
taining calcium  also.    In  fact,  limestone  (p.  169)  is  the  only 
rock  of  importance  which  is  not  made  up  of  silicon  com- 
pounds. 

275.  Properties. — Silicon  is  a  brown,  lusterless  powder 
which  melts  and  volatilizes  only  at  the  very  highest  tem- 
peratures.    It  is  unaffected  by  water  and  by  acids,  except 
hydrofluoric  acid,  HF,  which  produces  silicon  fluoride  and 

14  195 


196  ELEMENTARY  CHEMISTRY 

hydrogen.    When  heated  in  air  or  oxygen,  it  burns  brightly 
to  the  oxide. 

There  is  another  modification  of  the  element.  This 
forms  black  shining  crystals,  which  do  not  burn  when  heated 
in  the  air  or  in  oxygen,  and  are  not  acted  upon  by  hydro- 
fluoric acid. 

276.  Hydrogen  silicide,  SiH4. — Some  finely  powdered 
quartz  is  mixed  with  about  1  -J  times  its  weight  of  powdered 

magnesium  and  the  mixture  care- 
fully heated  in  a  dry  test-tube. 
It  is  sufficient  to  heat  one  point 
of  the  mass ;  this  starts  the  reac- 
tion, and  the  whole  contents  of 
the    tube    become    incandescent. 
When  the  tube  cools,  we  find  in 
FIG.  37. -Hydrogen  silicide.       ft   a   black,   friable   mass,   which 
is  mainly  magnesium  silicide,  SiMg2.     The  following  reac- 
tion has  occurred: 

SiQ2  +  4Mg  =  SiMg2  +  SMgO. 

When  some  of  the  powdered  substance  is  thrown  into 
dilute  hydrochloric  acid  in  a  beaker,  bubbles  of  a  colorless 
gas  rise  through  the  acid  and  take  fire  spontaneously  as 
soon  as  they  reach  the  surface,  burning  with  a  slight  ex- 
plosion and  giving  off  puffs  of  white  smoke.  This  gas  is 
hydrogen  silicide,  SiH4,  produced  thus : 

SiMg2  +  4HC1  =  SiH4  +  2MgCl2. 

In  preparing  larger  quantities  of  the  gas,  the  apparatus 
shown  in  Fig.  37  is  used.  The  magnesium  silicide  is  first 
introduced  into  the  gas-generating  bottle,  which  is  small. 
Then  water  is  poured  in  through  the  funnel-tube  until  all 
air  has  been  expelled.  This  is  necessary,  for  if  any  air  were 
left,  the  first  SiH4  evolved  would  ignite  spontaneously  and 
produce  an  explosion. 


THE  CARBON  GROUP  197 

The  next  step  is  to  pour  strong  hydrochloric  acid  gently 
through  the  funnel-tube.  A  colorless  gas  escapes  in  bubbles 
through  the  water  in  the  dish.  Each  bubble  burns  spon- 
taneously, producing  a  bright  flash  of  light  and  a  white  ring 
of  smoke.  This  smoke  is  silicon  oxide,  Si02;  the  burning 
of  the  gas  is  described  by  this  equation: 

SiH4  +  202  =  Si02  +  2H20. 

There  is  nothing  mysterious  about  this  spontaneous  in- 
flammation. Iron  burns  only  at  very  high  temperatures; 
charcoal  at  a  red  heat;  illuminating  gas  catches  fire  from 
a  flame  but  not  from  a  match-stick  bearing  a  spark;  the 
\apor  of  carbon  disulphide  is  inflamed  by  a  glass  rod  heated 
slightly  above  100°;  finally,  the  temperature  of  ignition  of 
the  gas  we  have  just  made  is  below  the  ordinary  tempera- 
ture of  the  air.  No  doubt  it  would  be  possible  to  cool  our 
gas  and  air  separately  to  such  a  temperature  that  no  com- 
bustion would  occur  when  they  were  brought  together, 
though  this  experiment  has  never  teen  tried. 

Pure  hydrogen  silicide  does  not  take  fire  spontaneously, 
though  it  is  very  inflammable.  Our  gas  contains  a  little 
hydrogen,  and  this  mixture  possesses  a  lower  temperature 
of  ignition  than  the  pure  gas. 

277.  Silicon  oxide,  Si02 — also  called  silica — results  when 
silicon  is  burned  in  the  air  or  when  silicon  hydroxide  is 
heated.  Obtained  in  this  way,  it  is  a  loose  white  powder, 
insoluble  in  water  and  in  acids  except  l^drofluoric  acid, 
which  converts  it  into  silicon  fluoride  and  water: 

Si02  +  4HF  =  SiF4  +  2H20. 

Silicon  oxide  melts  to  a  thin  liquid  in  the  oxyhydrogen 
flame  and  vaporizes  at  the  temperature  of  the  electric  arc. 
It  occurs  in  nature  in  great  abundance,  as  the  mineral 
quartz,  which  forms  crystals  like  those  shown  in  Fig.  38. 
Pure  quartz  is  perfectly  colorless  and  transparent,  and 
owing  to  this  it  was  supposed,  until  very  recent  times 


198 


ELEMENTARY   CHEMISTRY 


FIG.  38.— A  mass  of  quartz  crystals. 


of  the  sixteenth  century),  to  be  water  which  had  been  sub- 
jected to  intense  cold  and  so  completely  frozen  that  it  was 

impossible  to  melt 
it.  Quartz  is  often 
colored  by  traces 
of  impurities,  prob- 
ably compounds  of 
carbon  and  hydro- 
gen with  other  ele- 
ments.  Colored 
violet,  it  forms  the 
amethyst,  while 
transparent,  yellow 
quartz  is  a  cheap 
variety  of  topaz. 
Quartz  is  quite  hard,  and  is  very  inert  chemically.  It  is 
unaffected  by  the  agencies  which  dissolve  most  other  min- 
erals or  cause  them  to  crumble  and  form  soil.  Thus,  when 
a  rock  containing  it  is  destroyed,  the  quartz  remains.  In 
this  way  it  comes  about  that  the  sand  of  sea-  and  river- 
beaches  consists  very  largely  of  quartz.  When  such  sand 
becomes  cemented  to  a 'rock,  the  result  is  what  is  called  a 
sandstone.  Owing  to  the  hardness  of  quartz,  fine-grained 
sandstones  make  excellent  whetstones. 

Silica  also  occurs  in  nature  uncrystallized.  This  is 
called  opal.  Its  composition  is  the  same  as  that  of  quartz, 
but  it  is  more  impure,  frequently  containing  small  quan- 
tities of  iron  and  aluminium  oxides  and,  almost  invariably, 
water.  When  it  is  compact,  translucent,  and  exhibits  a 
beautiful  play  of  colors,  it  forms  the  gem  opal. 

278.  ~8rHxwn  hydroxide  (silicic  acid). — When  we  studied 
nitrous  acid  we  made  the  acquaintance  of  an  acid  which  is 
known  only  in  its  salts,  and  when  liberated  from  them  at 
once  decomposes  (p.  117).  Silicic  acid  behaves  in  the  same 
way.  Silicon  being  quadrivalent,  the  composition  of  its 


THE  CARBON  GROUP  199 

4     S  *  0 

hy4*axide — which  is  called  silicic  acid — would  be  Si(OH)4. 
Salts  in  which  its  hydrogen  is  replaced  by  metals  are  among 
the  most  important  minerals.  They  are  called  silicates. 
Thus,  magnesium  silicate,  Mg2Si04,  is  called  olivirie,  and  a 
silicate  in  which  half  of  the  hydrogen  is  replaced  by  copper 
is  called  dioptase,  CuH2Si04.  Hundreds  of  silicates  occur 
in  nature,  and  many  of  them  are  much  more  complex  in 
composition  than  these,  but  they  are  all  compounds  of 
metals  with  silicon  and  oxygen. 

279.  Glass  has  a  similar  chemical  make-up.  When  silica 
is  fused  with  calcium  carbonate  and  sodium  carbonate,  cal- 
cium silicate  and  sodium  silicate  are  formed,  both  liquid 
at  the  temperature  of  the  glass  furnace.  These  two  liquids 
mix  in  all  proportions  (like  water  and  alcohol),  and  when 
the  substance  cools  it  solidifies  without  any  separation, 
forming  a  solid  solution  of  sodium  and  calcium  silicate. 
If  the  glass  is  cooled  too  slowly  so  that  there  is  time  for 
crystallization  to  take  place,  the  solid  solution  is  trans- 
formed into  a  mass  of  small  crystals  of  sodium  silicate  and 
of  calcium  silicate;  it  becomes  white  and  opaque,  and  the 
glass  is  spoiled.  This  interesting  change  is  called  devitrifi- 
cation, and  it  takes  place  even  in  solid  cold  glass,  though 
very  much  more  slowly,  so  that  very  old  glass  frequently  be- 
comes dull  white  and  semi-opaque. 

Ordinary  window  and  bottle  glass  contains  mostly  so- 
dium and  calcium  silicates,  and  the  green  color  which  it 
often  has  is  due  to  the  presence  of  a  little  iron  silicate. 
When  lead  oxide  is  added  to  the  materials  in  the  melting- 
pot,  the  resulting  glass  will  also  contain  lead  silicate.  This 
gives  it  a  greater  density  and  a  higher  refractive  index. 
Such  glass  is  used  for  lenses  and  prisms. 

When  potassium  carbonate,  instead  of  sodium  carbo- 
nate, is  melted  with  quartz  and  calcium  carbonate,  the  glass 
will  contain  potassium  and  calcium  silicates.  Such  glass  is 
called  Bohemian  glass.  More  sand  is  used  in  making  it, 


200  ELEMENTARY  CHEMISTRY 

and  it  therefore  contains  a  higher  percentage  of  Si02  than 
ordinary  glass.  It  melts  at  a  higher  temperature  and  is  less 
attacked  by  most  liquids.  For  this  reason  it  is  largely  used 
in  making  chemical  glassware. 

It  will  be  seen  that  glass  varies  greatly  in  composition. 
This  agrees  with  the  statement  that  it  is  a  solid  solution, 
not  a  chemical  compound.  In  the  latter  case  such  variation 
would  be  impossible  (p.  33). 

280.  Sodium  silicate  is  soluble  in  water,  and  is  called 
water-glass.  We  should  expect,  by  adding  an  acid  to  this 
solution,  to  obtain  silicic  acid,  thus : 

Na4Si04  +  4HC1  =  4NaCl  +  Si(OH)4. 

When  the  experiment  is  tried,  a  colorless  jelly  separates 
which,  when  it  dries,  becomes  a  white  solid.  However,  anal- 
ysis shows  that  neither  the  jelly  nor  the  solid  is  silicic  acid, 
but  simply  a  mixture  of  Si02  with  variable  quantities  of 
water. 


CHAPTER   XXVIII 

THE  CARBON  GROUP  (Continued) 

TIN,  Sn  =  119. 

281.  Occurrence. — Bronze,  the  alloy  of  tin  and  copper, 
was  known  in  prehistoric  times,  and  tin  itself  was  known 
to  the  ancients,  who  considered  it  to  be  a  light-colored  and 
peculiar  variety  of  lead.     It  occurs  native  in  small  quanti- 
ties, but  the  most  important  ore  is  tin  dioxide,  Sn02,  called 
tin-stone  by  the  miners.    This  occurs  abundantly  in  south- 
western England  in  the  county  of  Cornwall,  in  Australia, 
and  especially  in  the  Malay  peninsula  and  the  neighboring 
islands. 

282.  Extraction. — The  first  step  is  to  free  the  mineral 
from  impurities  as  far  as  possible.    Sulphur  and  arsenic  can 
be  removed  by  roasting,  when  they  burn  away  as  oxides. 
Some  other  impurities  can  only  be  separated  by  a  process 
of  hand-picking.    Then  the  purified  tin-stone  is  heated  with 
charcoal  or  coal  in  suitable  furnaces.     The  tin  obtained  in 
this  way  is  purified  by  making  use  of  the  fact  that  its 
melting-point  (228°)  is  far  lower  than  that  of  the  impuri- 
ties it  contains.     Thus,  when  it  is  fused  at  a  low  tempera- 
ture, almost  pure  tin  runs  off  and  a  mass  containing  the 
iron,  copper,  and  other  metals  in  the  state  of  tin  alloys 
remains.     This  is  worked  up  again  to  remove  the  tin  it 
contains. 

283.  Properties. — Ordinary  tin  is  a  metal  with  a  bril- 
liant white  luster.    It  is  a  little  harder  than  lead  and  about 

201 


202  ELEMENTARY    CHEMISTRY 

seven  times  as  dense  as  water.  At  100°  it  is  ductile  and 
malleable;  at  200°  it  is  brittle  and  can  be  powdered.  It 
melts  at  a  low  temperature  and  boils  at  a  white  heat, 
catching  fire  if  in  contact  with  air,  and  burning  with  a 
bright  white  flame  to  tin  dioxide.  This  same  transforma- 
tion takes  place  more  slowly  at  a  lower  temperature,  when 
the  metal  is  melted  in  contact  with  air.  But  at  ordinary 
temperatures  it  is  unaffected  by  air  or  water. 

284.  Gray  tin. — There  is  a  second  modification  of  tin, 
a  gray,  crystalline  powder  much  less  dense  than  ordinary 
tin.     This  is  produced  when  tin  is  subjected  to  great  cold 
for  a  long  time — for  instance,  when  bars  of  tin  are  stored 
in  unheated  warehouses  during  cold  winters.     Sometimes 
the  tin  of  organ-pipes  in  unheated  churches  is  completely 
converted  into  the  gray  variety.    The  change  is  very  slow, 
but  may  be  made  much  more  rapid  by  placing  some  gray 
tin  in  contact  with  the  white  tin  and  putting  the  whole 
under  some  tin  chloride  solution  in  a  flask.     When  this  is 
done  it  is  found  that  below  20°  the  white  tin  turns  to  the 
gray  modification,  while  above  20°  the  change  takes  the  re- 
verse direction.     The  change  of  the  gray  variety  to  the 
white  is  rapid.     If  some  gray  tin  be  placed  in  a  beaker 
and  boiling  water  poured  over  it,  it  at  once  becomes  denser, 
acquires  a  white  metallic  luster,  and  changes  completely  to 
the  white  modification.    This  change  is  due  solely  to  the  rise 
in  temperature,  and  the  same  thing  happens  if  the  gray  tin 
is  heated  in  some  other  way. 

285.  Uses  of  tin. — Tin  is  used  in  the  manufacture  of 
solder,  an  alloy  of  tin  and  lead.     Tin-plate  is  sheet-iron 
coated  with  tin  by  cleaning  the  surface  carefully  and  dip- 
ping it  into  melted  tin.    It  is  employed  in  making  cans  for 
the  preservation  of  fruits.     The  acids  of  fruit  dissolve  a 
little  tin  from  such  vessels,  but  since  tin  compounds  in 
small  quantities  are  not  poisonous,  and  there  is  no  cumu- 
lative action  (pp.  146-207),  this  is  no  great  matter.    Some 


THE  CARBON  GROUP  203 

of  the  bar  tin  of  commerce,  especially  that  from  the  East,  is 
very  pure,  but  the  so-called  tin-foil,  which  is  used  for  wrap- 
ping food-products,  is  usually  highly  impure,  and  some- 
times is  nearly  pure  lead,  containing  no  tin  whatever. 

Tin  forms  two  series  of  compounds:  the  stannous  com- 
pounds, of  which  SnO  and  SnCl2  are  examples,  the  tin 
being  bivalent,  and  the  stannic  compounds — for  instance, 
Sn02  and  SnCl4 — in  which  the  tin  is  quadrivalent. 

286.  Stannous  chloride,  SnCl2,  is  the  most  important 
stannous  compound.    It  is  made  by  dissolving  the  metal  in 
hot,  strong  hydrochloric  acid.    When  the  solution  is  evap- 
orated it  separates  in  white  crystals  of  the  composition 
SnCl2.2H20,  which  are  called  tin-salt  by  the  dyer,  and  much 
used  as  a  mordant. 

The  stannous  compounds  tend  to  pass  into  stannic  com- 
pounds. Thus,  when  solutions  of  stannous  chloride  and 
mercuric  chloride  are  mixed,  then,  if  the  stannous  chloride 
is  in  excess,  a  grayish-black  precipitate  of  finely  divided  mer- 
cury is  produced,  stannic  chloride  being"  formed : 

HgCl2  +  SnCl2  =  SnCl4  +  Hg. 

Stannic  Compounds 

287.  Stannic  oxide,  Sn02,  is  found  in  nature  as  "  tin- 
stone," called  in  mineralogy  cassiterite.     Pure  tin-stone  is 
colorless,  but,  as  usually  found,  it  varies  from  yellow  to 
black.     Its  high  density  (7)  at  once  distinguishes  it  from 
most  other  minerals.    Stannic  oxide  may  be  made  artificially 
by  burning  tin  in  the  air.    It  then  forms  a  soft  white  pow- 
der, which  melts  with  difficulty  and  is  scarcely  affected  by 
acids. 

288.  Stannic  hydroxide,  Sn(OH)4,  is  an  interesting  sub- 
stance.   There  are  two  ways  of  making  it :  first,  by  precipi- 
tating a  solution  of  stannic  chloride,  SnCl4,  with  a  base : 

SnCl4  +  4KOH  =  Sn(OH)4  +  4KC1; 


204  ELEMENTARY  CHEMISTRY 

second,  by  the  action  of  strong  nitric  acid  on  tin.  The  two 
products  are  white  amorphous  powders  alike  in  appear- 
ance, and  have  exactly  the  same  composition,  but  they  are 
utterly  unlike  in  their  chemical  conduct.  Thus  the  hydrox- 
ide made  by  the  first  method  is  easily  soluble  in  the  three 
ordinary  acids,  but  that  prepared  by  the  second  is  not,  and 
this  difference  still  remains  when  the  two  products  have 
been  dissolved  separately  in  a  liquid  which  dissolves  them 
both,  and  then  obtained  again  in  the  solid  state. 

The  student  should  compare  this  curious  case  with  the 
two  modifications  of  mercuric  iodide  (p.  24).  The  state 
of  things  is  entirely  different.  The  red  and  the  yellow 
mercuric  iodide  melt  to  the  same  liquid,  yield  the  same 
vapor,  and  when  they  are  dissolved  separately  in  the  same 
solvent  (e.  g.,  alcohol)  the  two  solutions  produced  are  in 
all  respects  identical.  In  the  language  of  the  atomic  the- 
ory, we  may  say  that  the  two  are  different  only  in  the  way 
in  which  the  molecules  are  arranged  in  the  crystals.  Of 
course,  when  the  substance  is  liquefied,  vaporized,  or  dis- 
solved, the  crystals  vanish,  and  there  is  no  longer  any  ar- 
rangement. But  when  two  substances  remain  unlike  after 
such  treatment,  the  cause  of  the  difference  lies  deeper,  and 
can  only  be  explained  by  the  statement  that  the  atoms  in 
the  molecule  are  differently  arranged  in  the  two  cases. 

Substances  like  the  two  stannic  hydroxides,  which  are 
alike  in  composition  but  different  in  properties,  are  called 
isomeric. 

Stannic  hydroxide  in  both  its  forms  is  an  acid,  and 
readily  reacts  with  bases  forming  salts  in  which  its  hydro- 
gen is  replaced  by  metals.  It  is  often  called  stannic  acid 
for  this  reason,  and  the  salts  are  called  stannates. 

289.  Stannic  sulphide,  SnS2,  crystallizes  in  scales,  which 
have  the  color  and  something  of  the  luster  of  gold.  It  is 
called  mosaic  gold,  and  has  long  been  used  for  giving  a 
lustrous  metallic  surface  to  plaster  casts. 


THE  CARBON  GROUP  205 

Stannic  chloride,  SnCl4,  is  made  by  melting  tin  in  a 
retort  and  passing  a  current  of  dry  chlorine  over  it.  The 
vapor  is  led  through  a  condenser,  and  stannic  chloride  col- 
lects in  the  receiver  as  a  colorless,  fuming  liquid. 

Tin  amalgam  is  somewhat  employed  for  the  metallic 
reflecting  surface  with  which  the  glass  of  mirrors  is  backed, 
but  is  being  replaced  by  silver  for  this  purpose,  owing  to 
the  poisonous  action  of  the  mercury  upon  the  workmen. 


CHAPTEE   XXIX 

LEAD 
Pb  =  207. 

290.  Occurrence  and  extraction. — Lead  was  known  to 
the  ancients.    The  metal  itself  is  rare  in  nature.    The  car- 
bonate PbC03  and  the  sulphate  PbS04  occur  as  minerals, 
but  the  most  important  ore  is  galenite,  which  is  lead  sul- 
phide, PbS. 

In  order  to  obtain  the  lead  from  galenite,  it  is  roasted 
in  a  current  of  air,  the  sulphur  burning  away  as  sulphur  di- 
oxide, while  the  lead  remains  as  oxide.  The  lead  oxide  is 
then  heated  with  some  form  of  carbon  (charcoal  or  coal), 
when  carbon  dioxide  and  lead  are  produced. 

Another  method,  called  the  "air-reduction  process,"  is 
sometimes  used  with  very  pure  ores.  The  lead  sulphide  is 
roasted  until  a  portion  of  it  is  converted  into  lead  oxide; 
then  the  access  of  air  is  stopped  and  the  temperature  raised 
until  the  mixture  melts.  The  sulphur  of  the  lead  sulphide 
and  the  oxygen  of  the  lead  oxide  escape  together  as  sulphur 
dioxide,  while  lead  remain «  : 

2PbO  +  PbS  =  S02  +  3Pb. 

The  chief  lead-producing  countries  are  the  United 
States,  Spain,  and  Germany. 

291.  Properties. — Lead  is  softer  than  gold.    The  freshly 
cut  surface  has  a  bright  bluish-gray  metallic  luster,  but 
this  rapidly  disappears  by  oxidation,  giving  place  to  the 
familiar  dull  gray  color  of  the  metal.     This  attack,  how- 
ever, is  superficial,  and  the  metal  is  quite  permanent  in  air, 

206 


LEAD  207 

whether  moist  or  dry.  Lead  is  tolerably  malleable,  and  is 
plastic  at  a  gentle  heat,  so  that  it  can  be  made  into  tubes, 
but  it  is  not  tenacious  and  can  not  be  drawn  into  fine  wire. 

Lead  melts  readily  and  boils  at  a  white  heat.  Liquid  lead 
absorbs  oxygen  from  the  air,  passing  into  lead  oxide,  PbO. 

Pure  water  has  a  decided  action  on  lead,  converting  it 
into  the  hydroxide,  some  of  which  dissolves;  and  if  some 
fresh  shavings  of  the  metal  are  covered  with  distilled  water 
in  a  beaker  and  allowed  to  stand,  lead  can  be  detected  in 
the  liquid  after  half  an  hour.  This  fact  is  of  interest 
because  lead  pipes  are  used  for  conveying  water,  and  cases 
of  poisoning  have  been  brought  about  in  this  way.  How- 
ever, the  water  of  nature,  on  account  of  the  salts  dissolved 
in  it,  acts  upon  lead  much  less  rapidly  than  distilled  water 
does,  so  that  if  it  simply  runs  through  the  pipes  and  is  not 
allowed  to  stand  in  them,  there  is  little  danger. 

Lead  is  little  acted  upon  by  hydrochloric  or  sulphuric 
acid,  but  is  readily  dissolved  by  nitric  acid. 

Compounds  of  lead  are  poisonous.  It  requires  a  large 
quantity  of  a  lead  compound  to  produce  serious  symptoms 
if  it  is  all  introduced  into  the  system  at  once,  and  this  form 
of  lead-poisoning  rarely  occurs ;  but  chronic  lead-poisoning 
— that  is,  poisoning  from  the  continual  absorption  of  small 
quantities  of  lead  compounds  during  months— is  frequent 
and  dangerous.  All  workmen  who  deal  with  lead  and  its 
compounds  are  liable  to  it,  though  the  danger  can  be  re- 
duced by  perfect  cleanliness.  Some  of  the  symptoms  are 
violent  pains  in  the  abdominal  region,  a  loss  of  power  over 
po-H-Mn  muscles  (particularly  those  which  move  the  hand 
at  the  wrist),  and  a  slate-colored  line  round  the  edges  of 
the  frnms. 

292.  Compounds  of  lead  with  oxygen.— The  most  im- 
portant oxides  of  lead  are  lead  oxide,  PbO,  lead  dioxide, 
PbO*.  and  red  lead  or  minium,  "PKO^ 

Lead  oxide,  PbO,  is  made  by  heating  lead  above  its  melt- 


208  ELEMENTARY  CHEMISTRY 

ing-point  in  the  air.  It  occurs  in  commerce  as  a  heavy 
yellow  powder,  called  massicot,,  or  as  reddish-yellow  crys- 
talline scales,  called  litharge.  It  is  used  in  the  paint  manu- 
facture and  in  the  production  of  lead  glass  (p.  199)  It 
melts  easily,  and  when  heated  with  hydrogen  or  carbon  is 
converted  into  lead. 

Minium,  or  red  lead,  Pb304,  results  from  the  protracted 
heating  of  litharge  with  air-access.  It  is  a  heavy,  scarlet, 
crystalline  powder  which  is  used  as  a  color.  Its  formula 
may  be  written  thus : 

Pb022PbO. 

We  may  regard  it,  then,  as  a  compound  of  lead  oxide  and 
lead  dioxide.  When  it  is  treated  with  dilute  nitric  acid,  the 
lead  oxide  is  converted  into  lead  nitrate,  which  is  soluble 
and  can  be  washed  out  with  water,  while  lead  dioxide,  Pb(X, 
remains.  This  is  a  heavy,  dark-brown  powder,  insoluble 
in  water  and  in  nitric  acid.  It  yields  up  its  additional 
oxygen  readily,  and  when  a  little  of  it  is  ground  in  a  mor- 
tar with  some  sulphur,  the  mixture  is  inflamed.  Lead  diox- 
ide is  part  of  the  mixture  of  which  the  heads  of  matches 
consist. 

Both  lead  dioxide  and  red  lead  are  easily  converted  into 
lead  oxide  by  heat,  oxygen  escaping. 

Lead  is  quadrivalent  in  lead  dioxide,  and  in  lead  tetra- 
chloride,  PbCl4,  a  colorless,  very  unstable  liquid.  But  all  of 
its  important  compounds  correspond  to  lead  oxide,  and  in 
them  the  metal  is  bivalent.  Almost  all  lead  salts  are  in- 
soluble in  water,  or  nearly  so,  the  only  important  exceptions 
being  lead  nitrate  and  lead  acetate. 

293.  Lead  sulphide,  PbS,  is  the  mineral  galenite,  or  ga- 
lena. It  crystallizes  in  cubes,  which  have  a  lead-gray,  metal- 
lic luster  and  are  brittle.  It  is  produced  as  a  black  precipi- 
tate when  hydrogen  sulphide  is  passed  into  a  lead  solution : 

PbCL  +  H2S  =  PbS  +  2IIC1. 


LEAD  209 

The  insolubility  of  lead  sulphide  is  almost  complete,  and 
the  precipitate  forms  even  in  the  most  dilute  lead  solutions. 
This  is  therefore  a  very  delicate  test  for  lead. 

Lead  sulphate,  PbS04,  occurs  in  nature  as  the  mineral 
anglesite.  It  is  insoluble  in  water,  and  is  formed  as  a 

+4- 

heavy  white  precipitate  when  S04  ions  and  Pb  ions  come 
into  contact,  e.  g. : 

P"b(N08)a  +  H2S04  =  PbS04  +  2HN03. 

Lead  chloride,  PbCl2,  is  only  slightly  soluble  in  cold 
water  and  is  precipitated  when  chlorine  ions  and  lead  ions 
come  together,  unless  the  solution  is  dilute.  It  is  white, 
and  is  much  more  soluble  in  hot  water  than  in  cold. 

Lead  carbonate  is  white,  and  insoluble  in  water.  It 
occurs  in  nature  as  the  mineral  cerussite. 

A  substance  containing  lead  hydroxide  and  lead  carbon- 
ate in  about  the  proportions  expressed  by  the  following 
formula, 

Pb(OH)22PbC03, 

is  called  white  lead.  In  spite  of  its  active  poisonous  qual- 
ities, it  is  the  most  common  of  white  pigments,  and  is  the 
basis  of  most  ordinary  paint. 


NITROGEN  GROUP 


Nitrogen,  N. 
Phosphorus,  P. 
Arsenic,  As. 


Antimony,  Sb. 
Bismuth,  Bi. 


294.  We  have  here  an  example  of  the  fact  that,  in  the 
same  group,  the  metallic  properties  usually  increase  with 
the  atomic  weight.     Nitrogen   (N  =  14)   and  phosphorus 
(P  =  31)  are  non-metals,  arsenic  (As  =  75)  and  antimony 
(Sb  =  120)     are    metalloids,    bismuth    (Bi  =  208)     is    a 
metal.     Nitrogen,  phosphorus,  arsenic,  and  antimony  each 
form  a  gaseous  compound  with  hydrogen  of  the  formula 
RH3  (R  being  an  atom  of  the  element  in  question).     The 
stability  of  these  compounds  decreases  in  the  order  in  which 
the   elements   have   been  named,   with   increasing   atomic 
weight.    No  compound  of  bismuth  with  hydrogen  has  been 
obtained.    The  elements  of  this  group  are  trivalent  or  quin- 
quivalent.   Compounds  in  which  they  exhibit  other  valences 
are  known,  but  they  are  exceptional. 

PHOSPHORUS,  P  =  31. 

295.  History. — Phosphorus  was  discovered  in  the  latter 
part  of  the  seventeenth  century  by  Brand,  an  alchemist,  who 
was  attempting  to  prepare  a  liquid  which  would  turn  silver 
into  gold.    In  this  search  he  had  occasion  to  evaporate  urine 
to  dryness  and  heat  the  residue   strongly,   and   obtained 
phosphorus  among  the  products  which  condensed.     On  ac- 
count of  its  inflammability  and  its  luminosity  in  the  dark, 

210 


THE  NITROGEN  GROUP  211 

the  substance  excited  great  wonder,  and  was  exhibited  to 
various  princes,  among  others  to  Charles  II  of  England. 
Here  Boyle  made  its  acquaintance,  and  soon  after  worked 
out  for  himself  the  same  method  of  making  it;  but  the 
process  was  tedious  and  the  yield  very  small,  and  phos- 
phorus remained  a  curiosity  until,  about  a  century  after 
its  discovery,  Scheele  pointed  out  that  the  ashes  of  bones 
contained  it  in  abundance. 

296.  Occurrence, — Phosphorus  does  not  occur  free  in 
nature.    It  is  found  only  as  the  salts  of  its  most  important 
acid,  phosphoric  acid,  H3P04.    These  salts  are  called  phos- 
phates, and  calcium  phosphate,  Ca3(P04)2,  is   abundant. 
It  is  contained  in  many  minerals  and  rocks,  and  in  all 
fertile  soils.     From  the  soil  phosphorus  compounds  pass 
into  plants ;  cereals  especially  are  quite  rich  in  them.    The 
bones  and  teeth  of  mammals  are  largely  calcium  phosphate, 
and  brain   and  nerve  tissue   always   contain  considerable 
quantities  of  complex  phosphorus  compounds. 

297.  Preparation. — The  raw  material  for  the  prepara- 
tion of  phosphorus  is  calcium  phosphate,  Ca3(P04)2.    This 
is  used  either  as  bone-ash,  which  is  mainly  calcium  phos- 
phate, or  as  the  phosphate  rock,  which  occurs  naturally. 
It  is  mixed  with  charcoal  powder  and  fine  sand,  and  heated  to 
a  high  temperature  by  means  of  the  electric  arc  which  burns 
between  carbon  rods  projecting  into  the  mixture.     Phos- 
phorus is  liberated  as  vapor  which  is  condensed  under  water. 

298.  Properties. — Phosphorus  obtained  in  this  way  is 
a  colorless,  transparent  solid,  brittle  when  very  cold,  but 
at  ordinary  temperatures   soft  enough  to  be  cut  with  a 
knife.    It  is  insoluble  in  water  but  freely  soluble  in  carbon 
disulphide.     When  phosphorus  is  sealed  up  in  a  glass  tube 
from  which  the  air  has  been  removed,  and  allowed  to  stand 
in  a  dark  place,  it  slowly  vaporizes  and  condenses  on  the 
sides   of  the  tube  in  well-developed  transparent  crystals. 
The  phosphorus  of  commerce — usually  in  sticks  or  spheres 

15 


212  ELEMENTARY  CHEMISTRY 

— has  a  crystalline  structure,  but  a  confused  one,  no  crystal 
having  had  space  and  time  to  develop. 

Phosphorus  burns  brilliantly  to  phosphoric  oxide,  also 
called  phosphorus  pentoxide,  P205,  when  heated  gently  in 
the  air,  and  often  takes  fire  spontaneously.  At  ordinary 
temperatures  it  oxidizes  slowly,  and  this  oxidation  is  ac- 
companied by  the  production  of  ozone  and  by  the  evolution 
of  light  which  is  visible  in  the  dark.  Phosphorus  is  not 
luminous  when  placed  in  hydrogen  or  any  other  gas  which 
contains  no  free  oxygen. 

Nor  is  it  luminous  in  pure  oxygen.  If  a  stick  of  phosphorus  is 
transferred  in  the  dark  from  a  jar  of  air  (in  which,  of  course,  it  is 
giving  off  light)  into  a  jar  of  oxygen,  the  luminosity  is  quenched  at 
once.  But  the  phosphorus  begins  to  shine  again  when  some  of  the 
oxygen  in  the  jar  is  removed  by  an  air-pump  until  the  pressure  of 
that  which  is  left  is  about  £  of  an  atmosphere — in  other  words,  when 
the  jar  contains  only  about  the  same  quantity  of  oxygen  as  the  same 
volume  of  air.  In  this  way  it  can  be  shown  that  when  phosphorus 
is  placed  in  a  vacuum  and  oxygen  admitted,  it  begins  to  shine  as 
soon  as  the  vessel  contains  any  oxygen  at  all,  and  that  the  shining 
ceases  when  the  pressure  of  the  oxygen  has  reached  a  certain 
amount.  The  higher  the  temperature  the  greater  the  pressure 
required  to  stop  the  shining,  so  that  phosphorus  which  has  ceased 
to  be  luminous,  owing  to  too  much  oxygen  in  the  vessel,  again 
becomes  so  when  the  temperature  is  slightly  raised. 

299.  Effect   on   the   system. — Phosphorus   is   intensely 
poisonous,  and  on  account  of  the  ease  with  which  it  can  be 
obtained  from  the  heads  of  matches  is  frequently  made  use 
of  by  criminals.     There  is  also  a  chronic  phosphorus  poi- 
soning which  is  frequent  among  workmen  in  match-fac- 
tories, who  come  in  daily  contact  with  the  substance.    Two 
of  the  results  are  a  replacement  of  the  muscular  tissue  of 
certain  parts  of  the  body  by  fat  (notably  the  heart)  and  a 
distressing  decay  of  the  bones  of  the  jaw. 

300.  Bed  phosphorus. — This  is  a  dull  red  powder  of  im- 
perfect crystalline  structure.     It  is  not  luminous  in  the 


THE  NITROGEN  GROUP  213 

dark,  and  its  oxidation  in  the  air  is  immensely  slower  than 
that  of  the  colorless  variety.  Further,  the  temperature  at 
which  it  ignites  is  much  higher.  This  is  well  shown  by 
putting  a  small  quantity  of  each  modification  on  a  flat  brass 
rod,  a  few  inches  apart,  and  heating  the  rod  in  such  a  way 
that  the  red  phosphorus  is  nearest 
the  burner-flame  (Fig.  39).  The 
colorless  phosphorus  will  be  the  first 
to  inflame.  Eed  phosphorus  is  in- 
soluble in  carbon  disulphide  and  the 
other  liquids  which  dissolve  the  white 
variety.  Red  phosphorus  is  not  poi- 
sonous at  all  when  swallowed,  but  if 
it  is  suspended  in  water  and  the  liquid 
injected  into  a  vein,  phosphorus-poi-  p^.-Proof  that  red  phoe- 

SOning  results.  phorns  is  less  inflammable 

301.  Proof   that   red   and  white 
phosphorus  are  two  forms  of  the  same  element. — Why  do 

we  regard  two  bodies  so  totally  unlike  as  two  forms  of  the 
same  substance  ?  For  two  reasons  chiefly :  First,  because 
we  can  transform  either  completely  into  the  other  without 
adding  or  subtracting  anything  except  energy-;  second,  be- 
cause they  both  react  with  the  same  elements  to  form  the 
same  products.  Thus  both  burn  in  oxygen,  and  the  phos- 
phoric oxide  prepared  in  this  way  from  colorless  phosphorus 
is  identical  with  that  from  the  red  variety. 

Colorless  phosphorus  is  unstable  and  is  continually  pass- 
ing into  red  phosphorus,  which  is  the  natural  state  of  the 
element,  but  this  change  is  very  slow,  so  that  it  may  be 
kept  for  years  without  any  great  alteration.  Light  quick- 
ens the  change,  and  when  a  bottle  containing  sticks  of 
phosphorus  has  stood  for  a  long  time  in  the  same  position, 
that  side  of  each  stick  which  is  struck  by  the  light  becomes 
covered  with  a  film  of  red  phosphorus.  But  the  best  means 
of  accelerating  the  change  is  heat,  and  red  phosphorus  is 


214 


ELEMENTARY  CHEMISTRY 


made  by  heating  the  colorless  variety  to  260°  for  about  ten 
days  in  iron  vessels,  from  which  the  air  is  excluded  to 
prevent  combustion. 

302.  Uses  of  phosphorus. — The  chief  use  of  phosphorus 
is  for  the  production  of  matches.     The  match-sticks  are 
first  dipped  into  melted  paraffin  to  make  them  inflame  more 
easily,  and  then  into  a  paste  which  contains  phosphorus  with 
lead  dioxide,  Pb02,  or  some  other  substance  rich  in  oxy- 
gen, and  glue,  to  hold  the  mass  together.     They  are  then 
dried. 

303.  Compounds  of  phosphorus  and  hydrogen. — Phos- 
phine,  PH3,  is  made  by  heating  colorless  phosphorus  with 


FIG.  40.— Preparation  of  phoephine. 

ft  solution  of  sodium  hydroxide.  The  air  must  all  be  ex- 
pelled beforehand  from  the  retort  by  a  current  of  hydrogen 
or  illuminating  gas,  and  the  exit  tube  must  dip  under  warm 
water  (Fig.  40).  Colorless  bubbles  rise  through  the  water 
in  the  dish  and  inflame  spontaneously,  producing  white 
smoke-rings  of  phosphoric  acid,  H3P04: 

PH3  +  202  =  H3P04. 


, 

THE  NITROGEN  GROUP  215 

Pure  phosphine  is  a  colorless,  poisonous  gas  with  an 
odor  like  that  of  decaying  fish.  It  corresponds  in  composi- 
tion to  ammonia,  NH3  (p.  105),  and  we  should  expect  it, 
like  the  latter,  to  unite  with  acids  yielding  salts  corre- 
sponding to  the  ammonium  salts.  For  instance,  we  should 
expect  reactions  like  this — 

PH3  +  HC1  =  PH4C1— 

to  occur.  This  compound  PH4C1  corresponds  to  ammoni- 
um chloride,  and  is  called  phosphonium  chloride.  It  is  very 
unstable  and  can  only  exist  at  temperatures  far  below  0°. 
On  the  other  hand,  phosphonium  iodide,  PH4I,  is  well 
known  and  stable. 

Pure  PH3  does  not  inflame  spontaneously.  If  the  gas, 
made  as  described  above,  is  allowed,  before  escaping  into  the 
air,  to  pass  through  a  U-shaped  tube  immersed  in  a  freez- 
ing mixture,  there  condenses  in  this  tube  a  colorless  liquid, 
and  the  phosphine  which  escapes  no  longer  takes  fire  of 
itself,  though  it  is  still  very  inflammable.  Clearly  the  liquid 
which  collects  in  the  U-tube  must  be  the  cause  of  the  spon- 
taneous ignition,  and  this  we  can  prove  by  opening  the  tube 
cautiously,  when  the  liquid  inflames  and  continues  to  burn 
with  a  bright  flame,  producing  a  dense  white  smoke.  The 
liquid  which  collects  in  the  tube  is  another  compound  of 
hydrogen  and  phosphorus.  It  corresponds  to  hydrazine, 
and  has  the  composition  P2H4. 


CHAPTER   XXXI 

OXIDES  AND  ACIDS  OF  PHOSPHORUS-HALOGEN  COMPOUNDS 

304.  Phosphorous  oxide,  P406,  is  obtained  as  a  white 
volatile  powder,  with  a  garlic  odor,  by  burning  phosphorus 
in  a  limited  supply  of  air.    It  burns,  when  heated  in  air  or 
oxygen,  to 

Phosphoric  oxide,  also  called  phosphorus  pentoxide, 
P205.  This  is  the  product  of  the  combustion  of  the  element 
when  the  supply  of  oxygen  is  plentiful.  It  is  a  loose,  white, 
odorless  powder  with  an  acid  taste.  Its  attraction  for  water 
is  extraordinary,  and  the  best  means  of  drying  a  gas  is  to 
allow  it  to  stand  for  a  time  in  contact  with  this  sub- 
stance. 

305.  Acids  containing  phosphorus. — When  phosphorus 
pentoxide  is  thrown  into  water  it  dissolves  with  a  hissing 
noise,  and  the  liquid  becomes  hot.    Combination  has  taken 
place,    and    the    solution    contains    metaphosphoric    acid, 

HP03— 

P205  +  H20  =  2HP03. 

The  metaphosphoric  acid  can  be  obtained,  as  a  glassy 
mass,  by  evaporating  the  solution  to  dryness  and  heating 
the  residue.  But  if  the  solution  be  allowed  to  stand,  the 
metaphosphoric  acid  combines  with  more  water  and  passes 
into  phosphoric  acid,  H3P04 — 

HPO,  +  H20  =  H,P04. 

Phosphoric  acid  is  the  most  important  acid  of  phos- 
phorus. It  forms  hard,  brittle,  colorless  crystals,  very  solu- 


OXIDES  AND  ACIDS  OF  PHOSPHORUS  217 

ble  in  water.  The  salts  of  this  acid,  the  phosphates,  are  the 
only  compounds  of  phosphorus  found  in  nature.  Phos- 
phoric acid  contains  three  atoms  of  hydrogen  which  can  be 
replaced  by  metals,  and  therefore  each  metal  can  form 
three  phosphates,  e.  g.  : 

(1)  NaH2P04,  monosodium  phosphate; 

(2)  Na2HP04,  disodium  phosphate; 

(3)  Na3P04,  trisodium  phosphate; 

while  with  a  bivalent  metal,  like  calcium,  the  corresponding 
salts  will  be 

(1)  Ca(H2P04)2; 

(2)  CaHP04;and 

(3)  Ca3(P04)2. 

Salts  of  the  first  type  are  soluble  in  water.  Those  of 
the  second  and  third  are  insoluble  in  it,  except  with  potas- 
sium and  sodium,  all  of  whose  salts  are  soluble.  All  of  the 
phosphates  are  soluble  in  nitric  acid  except  tin  phosphate. 

A  hydroxide  in  which  phosphorus  is  quinquivalent  would  have 
the  formula  P(OH)6.  This  compound  is  unknown,  but  phosphoric 
acid  may  be  looked  upon  as  derived  from  it  by  the  loss  of  a  mole- 
cule of  water  ;  thus 

„    n     /O—  H  /O—  H 

5    n>p-°—  H-H.O  =  0=P-0—  H 

M)—  H  NO-H 

Phosphoric  acid. 

Metaphosphoric  acid  can  be  regarded  as  derived  from  the  same 
hydroxide  by  the  loss  of  two  molecules  of  water  : 


Metaphosphoric  acid. 

The  hydroxide  in  which  phosphorus  is  trivalent  is  called  phos- 
phorous acid,  P(OH)3.     It  can  be  obtained  by  treating  phosphorus 

trichloride,  PC1S,  with  water  — 

/O—  H 

PCls  +  3H2O  =  P-O-H  -f  3HC1. 
M)-H 


218 


ELEMENTARY  CHEMISTRY 


It  is  a  colorless,  crystalline,  deliquescent  solid,  which  tends  con- 
stantly to  absorb  more  oxygen  and  pass  into  phosphoric  acid. 

A  number  of  other  substances  containing  phosphorus, 
hydrogen,  and  oxygen  is  known.  These  substances  are  all 
acids,  and  are  more  or  less  similar  to  those  we  have  dis- 
cussed. 

306.  Compounds  of  phosphorus  and  sulphur. — Four  com- 
pounds of  these  two  elements  have  been  made  by  heating 


FIG.  41.— Preparation  of  phosphorus  trichloride. 

them  together  in  vessels  from  which  the  air  has  been  ex- 
pelled by  a  current  of  carbon  dioxide.  Eed  phosphorus  must 
be  employed,  for  explosions  occur  when  the  colorless  variety 
is  used.  The  compound  corresponding  to  phosphorus  pent- 
oxide  is  phosphorus  pentasulphide,  P2S5.  It  forms  yellow 
crystals,  which  melt  and  vaporize  without  decomposition  if 
air  is  excluded.  If  not,  they  take  fire  and  burn  with  a  pale 
flame.  This  substance  is  of  frequent  use  in  the  chemistry 


OXIDES  AND  ACIDS  OF  PHOSPHORUS  219 

of  the  carbon  compounds,  for  when  it  acts  upon  them  it  fre- 
quently removes  the  oxygen  and  inserts  sulphur  in  its  place. 
Thus,  when  alcohol,  C2H60,  is  treated  with  it,  the  corre- 
sponding sulphur  compound,  called  mercaptan,  C2H6S,  is 
obtained. 

The  other  sulphides  of  phosphorus  are  similar  to  phos- 
phorus pentasulphide. 

307.  Compounds  of  phosphorus  with  chlorine. — When 
phosphorus  is  melted  in  a  retort  and  a  current  of  chlorine 
passed  over  it,  it  burns  with  a  pale  flame,  and  phosphorus 
trichloride,  PC13,  passes  off  as  a  vapor,  which  can  be  con- 
densed by  cooling  (Fig.  41).  It  is  a  colorless  liquid  with 
a  pungent  odor.  The  vapor  attacks  the  eyes,  causing  tears. 
Its  behavior  with  water  has  just  been  discussed.  When  ex- 
posed to  air,  it  slowly  absorbs  oxygen,  passing  into  phos- 
phorus oxychloride,  POC13,  which  resembles  it  very  closely. 
When  more  chlorine  is  allowed  to  act  upon  phosphorus  tri- 
chloride, combination  takes  place  and  phosphorus  penta- 
chloride,  PC15,  is  produced  in  white,  lustrous  crystals. 


CHAPTER    XXXII 

THE  NITROGEN  GROUP  (Continued) 

ARSENIC,  As  =  75. 

308.  Occurrence. — Native  arsenic  is  rather  widely  dis- 
tributed.    Compounds  of  the  element  also  occur.     Impor- 
tant among  these  is  iron  sulph-arsenide,  FeSAs,  from  which 
arsenic  can  easily.be  obtained  by  heating  and  condensing 
the  vapor  which  is  given  off: 

FeSAs  =  FeS  +  As. 

Many  of  the  metallic  sulphides  which  are  important  as 
ores  contain  arsenic.  When  such  ores  are  roasted  to  produce 
the  oxides,  from  which  the  metals  are  afterward  to  be  ex- 
tracted, the  arsenic  burns  to  arsenious  oxide,  As406,  which  is 
a  frequent  by-product  in  such  operations.  Minute  traces  of 
arsenic  are  contained  in  animal  tissues — for  instance,  in 
the  human  thyroid  gland,  skin,  and  hair.  The  entire  human 
body  contains  about  half  a  milligram  of  the  element. 

309.  Properties. — Arsenic  exists  in  several  modifications. 
When  arsenic  vapor  is  suddenly  cooled  it  condenses  to  a  yel- 
low powder,  which  oxidizes  readily  in  the  air,  dissolves  in 
carbon  disulphide,  and  in  many  respects  exhibits  similarity 
to  colorless  phosphorus.     It  is  unstable  and  easily  passes 
into  crystallized  arsenic.     There  is  also  a  black  amorphous 
variety  which  is  much  more  stable  and  more  easily  obtained 
than  the  yellow.     Crystallized  arsenic  is  nearly  tin-white, 
with  a  metallic  luster,  brittle,  and  easily  powdered.    When 
heated  under  the  ordinary  pressure  of  the  air,  arsenic  vola- 


THE  NITROGEN  GROUP  221 

tilizes  without  melting,  but  it  may  be  melted  under  strong 
pressure.  The  vapor  is  lemon-yellow  and  poisonous.  Its 
density,  referred  to  hydrogen,  is  150.  Therefore,  the  molec- 
ular weight  must  be  150  X  2  =  300.  Accordingly,  the 
formula  of  arsenic  is  As4;  there  are  four  atoms  in  the 
molecule. 

310.  Arsine,  AsH3,  corresponds  to  ammonia.  It  can  be 
obtained  pure  by  treating  sodium  arsenide  (made  by  heating 
sodium  and  arsenic  together)  with  hydrochloric  acid: 

AsNa3  +  3HC1  =  3NaCl  +AsH3. 
Sodium  arsenide. 

Arsine  is  a  colorless,  unpleasant-smelling  gas  which  has 
been  condensed  to  a  transparent  liquid.  A  temperature  ap- 
proaching redness  separates  it  into  hydrogen  and  arsenic. 
It  is  combustible,  burning  to  water  and  arsenious  oxide. 
Arsine  is  frightfully  poisonous,  and  the  greatest  care  is 
required  in  working  with  it.  Even  traces  of  it  in  the  air 
cause  dizziness,  breathlessness,  and  fainting. 

When  a  solution  of  arsenious  oxide,  or  some  other  ar- 
senic compound,  is  added  to  a  liquid  in  which  hydrogen  is 
being  generated,  arsine  is  produced,  and  escapes  mixed  with 
the  hydrogen :  * 

As406  +  24H  =  4AsH3  +  6H20. 

1  On  the  other  hand,  if  the  hydrogen  is  generated  in  another  vessel 
and  passed  into  the  arsenic  solution  through  a  glass  tube,  there  is  no 
action.  Thus  it  will  be  seen  that  hydrogen,  in  the  moment  of  its  lib- 
eration, can  cause  a  chemical  change  which  free  hydrogen,  ready 
formed,  is  powerless  to  effect.  Here  is  another  example:  If  silver 
chloride  is  suspended  in  water,  and  hydrogen  passed  through  the 
water,  nothing  occurs;  but  if  zinc  and  sulphuric  acid,  or  any  other 
substances  which  can  generate  hydrogen,  are  added  to  the  same  liquid, 
the  silver  chloride  is  rapidly  converted  into  metal: 

AgCl  +  H  =  HC1  +  Ag. 

From  the  standpoint  of  the  atomic  theory,  such  behavior  is  explained 
by  the  statement  that  hydrogen  is  first  liberated  as  atoms.  Afterward 


222 


ELEMENTARY   CHEMISTRY 


This  operation  can  be  conducted  in  the  apparatus  repre- 
sented in  Fig.  42.  Zinc  and  dilute  sulphuric  acid  are  placed 
in  the  generator,  and  when  hydrogen  is  being  evolved  in 
a  regular  stream,  a  few  drops  of  a  water  solution  of  arseni- 


Fio.  42.— Marsh's  test. 

ous  oxide  are  added  through  the  funnel-tube.  The  colorless 
hydrogen  flame,  burning  at  the  jet,  enlarges,  becomes  vio- 
let-gray, and  gives  off  a  white  smoke  of  arsenious  oxide 
from  the  combustion  of  the  arsine.  If  the  flame  is  cooled 
by  placing  a  porcelain  dish  in  it,  arsenic  condenses  in 
blackish-brown  spots  on  the  dish. 

If  the  tube  through  which  the  mixture  of  hydrogen  and 
arsine  passes  is  heated  by  a  burner  flame,  the  arsine  is  de- 
composed by  the  heat  and  a  steel-gray  mirror  of  arsenic 
forms  in  the  tube.  The  formation  of  this  mirror  is  a  mar- 
velously  delicate  method  of  detecting  the  presence  of  arsenic, 
and  is  much  used  in  cases  of  suspected  poisoning.  It  is 
called,  after  its  discoverer,  Marsh's  test. 

these  unite  to  form  molecules,  H2,  and  then  they  become  less  ener- 
getic, because  the  bond  between  the  two  must  be  ruptured  before  any 
chemical  action  can  occur.  This  condition  of  activity  at  the  moment 
of  production  is  called  the  nascent  state,  and  many  similar  cases  are 
known,  not  only  with  hydrogen,  but  with  other  elements. 


THE  NITROGEN  GROUP  223 

311.  Arsenious  oxide,  As406,  occurs  in  commerce  as  a 
dense,   white   crystalline   powder.      By   slowly   cooling   its 
vapor  it  can  be  obtained  in  a  glassy,  amorphous  modifica- 
tion, which  becomes  opaque  and  white  on  being  preserved, 
passing  into  an  aggregate  of  small  crystals.1     Arsenious 
oxide  dissolves  very  slowly  in  water  in  the  cold,  but  consid- 
erable quantities  of  it  will  dissolve  if  time  enough  be  given. 
It  is  often  called  white  arsenic,  or  simply  arsenic.     It  is 
highly  poisonous,  and  owing  to  the  ease  with  which  it  can 
be  obtained  and  to  the  fact  that  its  taste  is  very  feeble,  it 
is  more  frequently  employed  by  poisoners  than  any  other 
substance.     The  poisons  of  the  Borgias  and  the  famous 
"  Acqua  Toffana,"  with  which,  it  is  said,  six  hundred  persons 
were  slain,  were  in  all  probability  preparations  containing 
arsenious  oxide. 

312.  Arsenic  pentoxide,  As205,  corresponding  to  P205, 
is  a  white  solid  which  dissolves  in  water,  producing  arsenic 
acid,  H3As04,  corresponding  to  phosphoric  acid,  H3P04. 
The  salts  of  arsenic  acid,  the  arsenates,  strongly  resemble 
the  phosphates,  and  usually  crystallize  in  the  same  forms. 
They  are  mostly  insoluble  in  water,  but  those  of  the  metals 
of  the  sodium  group  are  soluble. 

313.  Three   sulphides   of   arsenic   have   been   made   by 
melting  together  the  two  elements  in  the  proper  propor- 
tions. 

Arsenic  disulphide,  As2S2,  is  red,  crystalline,  and  trans- 
parent. It  was  formerly  used  as  a  paint  under  the  name 
realgar. 

Arsenic  trisulphide,  As2S8,  is  yellow,  and  is  still  occa- 
sionally used  as  a  color.  When  hydrogen  sulphide  is  passed 
into  a  solution  of  arsenious  oxide  in  water,  the  liquid  turns 

1  The  amorphous  arsenious  oxide  is  more  soluble  in  water  than  the 
crystalline.  In  fact,  different  modifications  of  the  same  solid  always 
have  different  solubilities  in  the  same  liquid,  the  uns'table  form  being 
the  more  soluble. 


224  ELEMENTARY  CHEMISTRY 

deep  yellow,  but  remains  clear.  Arsenic  trisulphide  has  been 
formed  thus : 

As406  +  6H2S  =  2As2S3  +  6H20; 

but  though  insoluble  in  water  when  once  formed,  it  does 
not  precipitate,  but  remains  in  the  liquid  in  a  curious  state 
intermediate  between  solution  and  suspension.  This  con- 
dition is  called  colloidal  solution.  It  is  different  from  a 
suspension  in  passing  through  a  filter  without  leaving  any 
solid  matter  upon  it,  and  in  being  apparently  clear.  Many 
other  colloidal  solutions  are  known.  They  behave  like  sus- 
pensions in  which  the  suspended  substance  is  very  finely 
divided. 

The  addition  of  a  few  drops  of  hydrochloric  acid  will 
cause  the  arsenic  trisulphide  to  separate  in  yellow  flakes,  and 
it  is  now  entirely  insoluble. 

Arsenic  pentasulpliide,  As2S5,  is  yellow  and  unstable, 
tending  to  separate  into  arsenic  trisulphide  and  sulphur. 


CHAPTER     XXXIII 

THE  NITROGEN   GROUP   (Continued) 

ANTIMONY  AND  BISMUTH 
ANTIMONY,  Sb  =  120. 

314.  Preparation. — Antimony  occurs  native,  but  its  most 
important  ore  is  the  sulphide  Sb2S3,  which  is  called  stibnite. 
From  this  the  element  is  obtained  by  strongly  heating  with 
iron  filings  in  a  well-covered  crucible — 

Sb2S3  +  3Fe  =  3FeS  +  2Sb— 

or  by  carefully  roasting  the  sulphide  in  a  furnace  until  it  is 
converted  into  antimony  oxide.  This  is  then  mixed  with 
powdered  coal,  or  some  other  form  of  carbon,  and  heated 
to  faint  redness,  when  carbon  dioxide  escapes  and  antimony 
remains. 

315.  Properties. — Antimony  is  a  crystalline,  brittle  sub- 
stance with  a  silver-white,  metallic  luster.    It  melts  easily 
and  volatilizes  at  a  strong  red  heat.     It  is  not  affected  by 
air  at  ordinary  temperatures,  but  when  heated  burns,  giving 
off  a  white  smoke  of  antimonious  oxide,  Sb406.    The  element 
is  employed  in  the  production  of  several  important  alloys. 
Type  metal  is  an  alloy  of  lead,  tin,  and  antimony ;  Britannia 
metal  an  alloy  of  antimony  and  tin. 

316.  Stibine,  SbH3,  resembles  arsine,  and  is  made  by 
similar  methods — e.  g.,  by  treating  an  alloy  of  zinc  and  anti- 
mony with  hydrochloric  acid : 

Sb2Zn3  +  6HC1  =  3ZnCl2  +  2SbHR. 

220 


226  ELEMENTARY  CHEMISTRY 

It  is  a  colorless  gas  with  an  unpleasant  odor,  less  poi- 
sonous than  arsine.  It  is  very  unstable,  separating  into  its 
constituents  even  at  ordinary  temperatures  and  rapidly 
when  heated.  For  this  reason  it  has  never  been  obtained 
pure,  and  our  knowledge  of  its  properties  has  been  obtained 
by  studying  a  mixture  of  the  gas  with  hydrogen.  If  a 
solution  of  an  antimony  compound  is  added  to  the  contents 
of  a  hydrogen  generator,  as  in  Marsh's  test  (p.  222),  stibine 
escapes,  mixed  with  the  hydrogen.  If  the  tube  through 
which  the  gas  passes  is  heated  the  stibine  is  decomposed  and 
an  antimony  mirror  is  deposited.  Or,  if  a  cold  porcelain 
dish  is  held  in  the  flame  of  the  burning  gas,  antimony  spots 
are  formed  upon  it.  These  deposits  of  antimony  are  not 
affected  by  a  solution  of  sodium  hypochlorite,  NaCIO,  which 
rapidly  dissolves  the  arsenic  mirror  and  spots,  and  this 
fact  enables  us  to  decide  which  element  is  present. 

317:  We  have  pointed  out  that  in  the  same  group  the 
elements  usually  become  more  metallic  with  increasing 
atomic  weight.  We  should  therefore  expect  antimony  to  be 
more  metallic  than  arsenic.  That  this  is  the  case  is  shown 
by  the  fact  that  antimony  is  able,  to  a  small  extent,  to 
replace  the  hydrogen  of  acids  forming  salts.  Arsenic  is 
not.  Thus,  antimony  sulphate,  Sb2(S04)3,  and  antimony 
nitrate,  Sb(N03)3,  are  known,  but  they  are  very  unstable. 

The  formula  of  stibine,  SbH3,  shows  that  antimony  is 
trivalent  toward  hydrogen.  The  valence  of  an  element 
toward  hydrogen  is  always  the  same.  Thus,  if  a  compound 
SbH5  were  obtained,  we  should  have  to  admit  that  anti- 
mony could  have  a  valence  of  five  as  well  as  three  toward 
hydrogen.  Such  cases  are  unknown.  However,  antimony 
is  both  trivalent  and  quinquivalent  toward  the  halogens, 
and  almost  all  the  compounds  SbX3  and  SbX5  (where  X 
is  a  halogen  atom)  are  known. 

Antimony  trisulphide,  Sb2S3,  is  the  mineral  stibnite.  It 
occurs  in  brittle,  heavy,  blackish-gray  crystals,  which  melt 


JUSTUS  VON  LIEBIG 
B.  Germany,  1803.     D.  Munich,  1873. 


THE  NITROGEN  GROUP  227 

easily,  even  in  the  flame  of  a  candle.  The  antimony  trisul- 
phide  obtained  by  precipitating  an  antimony  solution  with 
hydrogen  sulphide — 

2SbCl3  +  3H2S  =  Sb2S3  +  6HC1— 

is  an  orange-red,  amorphous  powder,  totally  different  from 
stibnite,  but  having  the  same  composition.  It  can  be  ob- 
tained gray,  metallic,  and  crystalline,  like  stibnite,  by  melt- 
ing it  and  letting  it  cool  slowly. 

BISMUTH,  Bi  =  208. 

318.  Bismuth  occurs  in  nature  chiefly  as  metal.  Com- 
pounds of  the  metal  are  found  also,  but  only  the  sulphide, 
"  bismuth  glance/7  Bi2S3,  possesses  any  importance  as  an 
ore.  The  usual  ores  contain  bismuth,  more  or  less  bismuth 
sulphide,  and  other  substances.  They  are  first  roasted, 
which  burns  away  the  arsenic  which  is  usually  present,  and 
converts  much  of  the  bismuth  sulphide  into  bismuth  triox- 
ide,  Bi203.  The  roasted  ore  consists,  therefore,  of  a  mixture 
of  bismuth,  bismuth  trioxide,  and  unaltered  bismuth  sul- 
phide, together  with  the  earthy  or  stony  impurities  called 
by  the  .miners  "  gangue."  It  is  mixed  with  carbon  and  iron 
filings  and  melted  in  crucibles.  The  carbon  removes  the 
oxygen  from  the  bismuth  oxide,  and  the  iron  the  sulphur 
from  the  bismuth  sulphide,  producing  in  both  cases  bismuth, 
which  mixes  with  that  which  was  present  at  the  beginning, 
and  which,  of  course,  is  melted  by  the  heat.  The  metal  col- 
lects in  a  layer  on  the  bottom  of  the  crucible. 

Bismuth  has  a  strong,  reddish-white  metallic  luster  and 
forms  crystals  which  look  like  cubes  but  are  not,  because 
measurement  shows  the  angles  to  be  oblique.  It  melts  very 
readily  (270°),  and  its  alloys  melt  at  still  lower  tempera- 
tures. An  intimate  mixture  of  two  substances  in  the  proper 
proportions  always  melts  at  a  lower  temperature  than  either 
alone,  provided  that  on  being  melted  they  form  a  solution. 
16 


•228  ELEMENTARY   CHEMISTRY 

Thus,  the  melting-point  of  ice  is  not  affected  by  mixing  it 
with  sand,  which  does  not  dissolve  in  the  water  produced 
by  the  melting,  but  when  the  ice  is  mixed  with  salt,  which 
does  dissolve,  the  ice  melts  at  a  lower  temperature.  Metals 
behave  in  the  same  way,  and  bismuth  is  used,  along  with 
cadmium,  tin,  and  lead  in  the  manufacture  of  fusible  alloys, 
some  of  which  melt  easily  in  hot  water  (p.  179). 

Bismuth  burns  to  Bi203,  when  heated  in  the  air.  Hy- 
drochloric acid  does  not  affect  it.  Hot,  strong  sulphuric  acid 
converts  it  into  sulphate.  Nitric  acid,  either  cold  or  hot, 
dissolves  it  easily  to  bismuth  nitrate,  Bi(N03)3. 

Bismuth  is  usually  trivalent.  No  compounds  are  known 
in  which  one  atom  of  bismuth  combines  with  five  halogen 
atoms. 

Bismuth  nitrate,  Bi(NOz)3,  results  when  the  metal  is 
dissolved  in  nitric  acid.  When  the  liquid  is  evaporated  it 
forms  colorless  crystals  which  contain  five  molecules  of 
water.  Like  other  bismuth  salts,  the  nitrate  is  decomposed 
by  water, 

Bi(N03)3  +  H20  =  BiON03  +  2HN03, 

Bismuth  oxynitrate 

bismuth  oxynitrate  being  deposited  as  a  white  powder  in- 
soluble in  water.  This  powder  is  largely  used  in  medicine 
under  the  name  "  subnitrate  of  bismuth." 

Bismuth  has  a  remarkable  tendency  to  form  compounds, 
like  BiOCl  and  BiON03,  which  are  partly  salts  and  partly 
oxides.  Such  compounds  are  called  "  basic  salts"  Some- 
times a  basic  salt  is  written  simply  as  a  compound  of  the 
salt  itself  with  the  oxide.  Thus,  basic  lead  chloride  may  be 
written  PbCl2.PbO.  A  compound  which  is  partly  salt  and 
partly  hydroxide  is  also  called  a  basic  salt.  Thus,  tin  basic 

Cl 
chloride  is  Sn 


CHAPTER    XXXIV 

THE  CHROMIUM  GROUP 

Chromium,  Cr.  Uranium,  U. 

CHROMIUM,  Cr  =  52. 

319.  Chromium  and  uranium  are  metals,  but  they  have 
one  important  non-metallic  characteristic — they  form  part 
of  the  make-up  of  acids.     Thus,  chromium  forms  an  acid, 
H2Cr04  (known  chiefly  in  its  salts),  and  similar  compounds 
of  uranium  are  known.    H2Cr04  is  called  chromic  acid,  and 
its  salts,  the  chromates,  exhibit  many  similarities  with  the 
sulphates. 

320.  Occurrence  and  preparation. — Chromium  has  not 
been  found  free  in  nature.    Its  most  abundant  compound  is 
chromite,  or  chrome  iron  ore,  which  may  be  regarded  as  a 
double  oxide  of  iron  and  chromium,  Cr203.FeO.    This  is  the 
source  of  the  chromium  compounds  of  commerce. 

For  a  century  after  the  discovery  of  the  element  (1797) 
the  preparation  of  chromium  was  difficult,  but  within  the 
last  few  years  two  excellent  methods  of  making  it  have 
been  devised  which  bid  fair  to  become  of  commercial  im- 
portance. 

First,  when  chromic  oxide,  Cr203,  is  heated  with  carbon 
to  the  temperature  of  the  electric  arc,  carbon  monoxide  es- 
capes and  chromium  is  produced. 

Second,  when  a  mixture  of  chromic  oxide  and  powdered 
aluminium  is  strongly  heated  at  one  point  the  whole  sub- 

229 


230  ELEMENTARY  CHEMISTRY 

stance  becomes  intensely  white-hot  and  is  converted  into 
aluminium  oxide  and  chromium : 

Cr203  +  2A1  =  A1203  +  2Cr. 

321.  Properties. — Chromium  is  a  metal  with  a  grayish- 
white  luster  which  is  only  slightly  tarnished  in  the  air.    It 
is  soft  enough  to  be  easily  filed,  and  is  not  attracted  by  the 
magnet.     It  does  not  melt  in  the  oxyhydrogen  flame.     At 
a  high  temperature  it  burns  in  oxygen  to  Cr203,  producing 
sparks  like  those  of  burning  iron,  but  more  brilliant. 

322.  Peculiar  behavior   toward    acids.— Hydrochloric   acid 
dissolves  chromium,  liberating  hydrogen,  and  it  frequently  happens 
that,  when  the  chromium  is  first  dropped  into  the  acid,  the  action  is 
very  slow,  but  gradually  gathers  strength  until  it  is  quite  energetic, 
then  diminishes  until  scarcely  any  hydrogen  is  liberated,  then  again 
increases,    and  so   on   until  the    chromium   is   all   dissolved.      If 
chromium  is  left  exposed  to  the  air  it  gets  into  what  is  called  the 
passive  state,  in  which  it  behaves  like  gold  and  is  insoluble  in  acids. 
But  if  the  chromium,  while  under  the  acid,  is  touched  with  a  piece 
of  zinc,  the  passive  state  is  destroyed,  hydrogen  is  liberated,  and  the 
chromium  suddenly  begins  to  dissolve. 

323.  Chromic  oxide,  Cr203,  can  be  obtained  by  heating 
the  hydroxide.    When  amorphous  it  is  bright  green,  when 
crystallized  blackish-green.    It  is  insoluble  in  acids  and  not 
affected  by  any  ordinary  heat.     Under  the  name  "  chrome 
green"  it  is  used  for  painting  on  porcelain. 

When,  in  a  series  of  salts,  a  metal  has  the  same  valence 
that  it  has  in  a  certain  oxide,  we  are  in  the  habit  of  saying 
that  the  salts  correspond  to  or  are  derived  from  that  oxide. 
Thus  the  common  salts  of  lead,  like  PbCl2  and  Pb(N03)2— 
in  which  the  lead  is  bivalent — correspond  to  litharge  PbO, 
and  not  to  Pb02  or  Pb304.  In  this  sense  the  ordinary  salts 
of  chromium — the  chromic  salts — are  derived  from  chromic 
oxide,  for  in  both  the  oxide  and  the  salts  the  metal  is  tri- 
valent. 

324.  Chromic  chloride,  CrCl3,  forms  green  crystals  with 


THE  CHROMIUM  GROUP  231 

6H20,  which  are  quite  soluble  in  water.  Anhydrous  CrCl3 
forms  violet-red  scales,  which  are  practically  insoluble  in 
water — a  remarkable  fact,  for  we  should  expect  the  scales 
to  combine  with  water,  producing  CrCl36H20,  and  then 
dissolve.1 

The  other  chromic  salts,  the  sulphate,  nitrate,  and  so  on, 
exhibit  the  same  curious  insolubility  of  the  anhydrous  salt. 

325.  Chromium  trioxide,  Cr03,  is  made  by  decomposing 
a  solution  of  potassium  dichromate,  K2Cr207,  with  sulphuric 
acid: 

K2Cr207  +  H2S04  =  K2S04  -f-H20  +  2Cr03. 

It  crystallizes  in  red  needles,  which  are  deliquescent  and 
very  soluble  in  water.  When  heated  gently  it  melts  without 
decomposition,  but  at  a  higher  temperature  it  separates  into 
chromic  oxide  and  oxygen.  Formerly  it  was  called  by  the 
incorrect  name  of  "  chromic  acid''  and  this  name  is  still 
in  use  in  commerce. 

326.  Chromic  acid,  H2Cr04,  is  obtained  in  small  rose- 
red  crystals  by  cooling  a  saturated  solution  of  chromium 
trioxide  in  water.    It  is  very  unstable,  easily  separating  into 
Cr03  and  H20.     Its  salts,  on  the  other  hand,  are  well 
known  and  stable.    They  are  called  the  chromates.    Chromic 
acid  corresponds  to  sulphuric  acid  H2S04,  and  the  chromates 
correspond  to  the  sulphates.    The  sulphate  and  the  chromate 
of  the  same  metal  usually  crystallize  in  the  same  form. 

Potassium  chromate,  K2Cr04,  forms  yellow  crystals, 
soluble  in  water  to  a  yellow  liquid. 

Lead  chromate,  PbCr04,  is  made  by  adding  a  solution 
of  potassium  chromate  (or  of  potassium  dichromate)  to  a 
solution  of  a  lead  salt,  thus : 

Pb(N03)2  +  K2Cr04  =  PbCr04  +  2KN03. 

1  The  explanation  is  that  the  speed  with  which  anhydrous  CrCl, 
dissolves  is  very  small.  In  presence  of  certain  substances,  which  act 
catalytically,  it  dissolves  rapidly. 


232  ELEMENTARY  CHEMISTRY 

Most  of  the  chromates  of  the  metals  are  yellow  or  red. 
Many  of  them  are  insoluble  in  water. 

327.  Potassium    dichromate,    K2Cr207,    is    made    from 
chrome  iron  ore  by  various  methods,  one  of  which  is  to  heat 
the  powdered  mineral  with  potassium  carbonate  in  a  current 
of  air.    It  serves  as  the  starting-point  for  the  preparation  of 
other  chromium  compounds.    At  present  the  chromate  and 
dichromate  of  sodium  are  replacing  the  potassium  salts  in 
commerce,  because  they  are  cheaper. 

The  dichromates  are  mostly  red,  and  soluble  in  water. 

Their  composition  can  be  remembered  by  imagining  them  to 
consist  of  the  chromate  with  OO3.  Thus  potassium  dichromate 
can  be  written  KaCrC^CrOs.  There  is  also  a  potassium  trichromate, 
KaCrO42CrO8,  and  a  potassium  tetrachromate,  KaCrO43CrO8. 

All  of  the  chromium  compounds  are  poisonous.  The 
chromates  and  dichromates  are  more  poisonous  than  the 
salts  in  which  chromium  acts  like  a  metal. 

328.  Uranium  is  comparatively  rare.     It  has  a  special 
interest  because  it  has  the  highest  atomic  weight  of  all  the 
elements  known  at  present    (U  =  239).      It  is  found  in 
nature  chiefly  as  the  mineral  pitchblende,  which  is  an  oxide 
of  the  composition  U308 — usually  very  impure.     The  best 
method  of  making  uranium  is  to  heat  purified  U308  with 
pure  charcoal  to  the  temperature  of  the  electric  arc.     It 
is  a  metallic  substance  with  a  pure  white  luster,  non-mag- 
netic, and  soft  enough  to  be  filed. 

When  a  piece  of  uranium  or  of  pitchblende  is  placed  in  the  dark 
near  a  photographic  plate,  and  the  plate  is  afterward  developed,  it 
is  found  to  be  affected  in  somewhat  the  same  way  as  if  the  uranium 
had  been  a  source  of  light  during  the  time  of  exposure.  Uranium 
and  its  compounds,  then,  have  the  power  of  sending  forth  rays  which 
affect  the  photographic  plate.  But  these  rays  differ  from  light  in 
many  remarkable  ways.  Thus,  they  pass  readily  through  black  paper 
and  through  thin  plates  of  aluminium  or  copper.  A  thick  plate  of 
platinum  or,  best  of  all,  a  plate  of  lead  will  stop  them.  Again,  the 
rays  pass  through  a  glass  prism  without  being  bent  out  of  their 


THE   CHROMIUM   GROUP  233 

course,  and  are  not  reflected  from  a  polished  metallic  surface  in  the 
same  way  as  light.  It  has  been  shown  that  these  curious  rays  are 
not  due  to  the  uranium  itself,  but  to  other  elements  which  are  con- 
tained in  the  pitchblende  in  small  quantities.  Radium  (p.  171)  is 
the  most  important  of  these  elements. 

The  question  how  the  radium  compounds  obtain  the  energy 
which  they  continuously  give  out  is  an  unsolved  riddle.  A  possible 
explanation  is  that  the  radium  atom  is  not  a  single  particle,  but  a 
group  of  numerous  smaller  particles,  all  in  rapid  motion.  When 
the  motion  of  a  particle  or  group  of  particles  becomes  very  energetic 
it  separates  from  the  rest  of  the  atom  and  shoots  out  with  great 
rapidity,  producing  the  rays.  At  the  same  time  the  shock  produces 
vibrations  of  the  ether  and  the  surrounding  matter  which  appear  as 
light  and  heat.  What  becomes  of  the  rest  of  the  radium  atom  after 
the  fragment  has  been  pitched  off  ?  It  can  no  longer  be  radium. 
There  is  some  evidence  that  the  other  product  is  helium — at  least  it 
appears  that  helium  is  constantly  produced  by  radium  compounds. 
This  would  realize  the  dream  of  the  alchemists — the  conversion  of 
one  element  into  another.  But  so  far  the  evidence  is  not  strong 
enough  to  establish  such  a  revolutionary  conclusion. 


CHAPTER    XXXV 

THE    OXYGEN    GROUP 
Oxygen,  O.  Sulphur,  S.          Selenium,  Se.          Tellurium,  Te. 

329.  The  elements  of  this  group  are  non-metals.    They 
all  exist  in  several  allotropic  modifications.     They  combine 
with  two  atoms  of  hydrogen,  and  the  products — except  in 
the  case  of  water,  H20 — are  colorless  gases  with  a  foul 
odor  and  intense  poisonous  action.     Oxygen  has  already 
been  discussed.     Sulphur  is  common.     Selenium  and  tellu- 
rium are  rare. 

SULPHUR,  S  =  32. 

330.  Occurrence   and   preparation. — Large   deposits   of 
native  sulphur  occur  in  volcanic  regions,  particularly  in 
Sicily,  and  most  of  the  sulphur  of  commerce  is  obtained 
from  this  source.     We  have  already  seen  that  many  sul- 
phides— for  instance,  ZnS  and  PbS — are  very  important  as 
ores  of  the  metals.     When  such  ores  are  roasted  in  order 
to  convert  the  metal  into  oxide  (afterward  to  be  heated  with 
carbon),  the  sulphur  burns  to  sulphur  dioxide,  which  is 
then  converted  into  sulphuric  acid  and  sold  in  this  form. 

A  little  of  the  sulphur  of  Sicily  occurs  in  pure  trans- 
parent yellow  crystals,  but  the  greater  portion  contains 
much  calcium  sulphate  and  earthy  matter.  The  crude  ore 
is  piled  in  large  heaps  in  shallow  pits  lined  with  plaster. 
These  heaps  are  covered  with  powdered  ore,  and  air-canals 
are  left  in  the  interior,  so  that  the  air  has  a  limited  access. 
Then  the  sulphur  in  the  lower  part  of  the  air-canals  is  ig- 
nited, and  a  smothered  combustion  lasting  some  weeks 
ensues,  during  which  nearly  half  of  the  sulphur  is  burned, 
234 


THE  OXYGEN    GROUP  235 

and  the  rest  is  melted  by  the  heat  and  flows  out.  This 
burning  of  part  of  the  sulphur  in  order  to  obtain  the  rest 
seems  like  a  wasteful  proceeding,  but  it  must  be  remem- 
bered that  fuel  is  scarce  in  Sicily,  and  the  sulphur  ore  is 
the  cheapest  combustible  at  hand. 

The  crude  sulphur  obtained  in  this  way  still  contains 
calcium  sulphate,  arsenic,  and  other  substances.  It  is  puri- 
fied by  distillation  from  cast-iron  retorts,  the  vapor  being 
led  into  a  stone  chamber,  on  the  floor  of  which  the  melted 
sulphur  collects.  From  time  to  time  it  is  withdrawn  and 
poured  into  cylindrical  moulds  of  wood,  in  which  it  solidifies 
to  form  roll  sulphur. 

Flowers  of  sulphur  is  obtained  by  leading  the  sulphur- 
vapor  into  larger  stone  chambers,  the  walls  of  which  remain 
cold.  The  sudden  cooling  condenses  the  vapor  to  a  fine 
yellow  powder,  which  is  at  first  amorphous,  but  gradually 
becomes  crystalline  on  being  kept. 

331.  Properties, — Sulphur  is  a  brittle  yellow  solid,  and 
is  practically  a  non-conductor  of  electricity  and  a  very  bad 
conductor  of  heat.  It  melts  slightly  above  100°  to  a  thin, 
light  yellow  liquid.  When  this  liquid  is  heated  to  a  higher 
temperature  it  forms  a  jelly-like  mass,  darker  in  color,  and 
so  thick  that  the  flask  containing  it  can  be  inverted  with- 
out the  contents  running  down  the  walls.  This  thickening 
is  accompanied  by  a  distinct  fall  in  temperature,  which 
shows  that  another  modification  of  sulphur,  richer  in  en- 
ergy, is  formed.  Heated  still  hotter,  it  again  becomes  a 
thin  liquid,  but  is  now  reddish  black  and  nearly  opaque. 
Finally  it  boils.  The  vapor  is  at  first  yellow,  but  its  color 
deepens  with  rising  temperature  until,  at  a  faint  red  heat, 
it  is  blood-red.  This  variation  in  color  would  lead  us  to 
suspect  that  some  alteration  in  the  sulphur  molecule  occurs 
on  heating.  As  a  matter  of  fact,  the  density  of  the  vapor 
shows  that  at  low  temperatures  the  formula  is  S8,  at  high 
temperatures  S2.  Sulphur  is  very  soluble  in  carbon  disul- 


236 


ELEMENTARY  CHEMISTRY 


FIG.  43.— Crystals  of  native 
sulphur. 


phide,  CS2,  and  in  some  other  liquids,  and  it  has  been  shown 
that  the  molecule  of  dissolved  sulphur  also  contains  eight 
atoms. 

Heated  in  the  air,  sulphur  takes  fire  and  burns  with  a 
pale  blue  flame,  mainly  to  sulphur  dioxide,  S02,  but  traces 
of  sulphur  trioxide,  S03,  are  produced  at  the  same  time. 
Even  at  ordinary  temperatures,  sulphur  oxidizes  slowly  in 
the  air,  and  if  some  powdered  sulphur  is  stirred  up  with 
water  for  a  time  and  the  mass  poured  upon  a  filter,  sul- 
phuric acid  can  be  detected 
in  the  clear  liquid  which  runs 
through : 

S  +  H20  +  30  =  H2S04. 

332.  Allotropic    modifications. 

— Sulphur  exists  in  a  number  of 
forms,  most  of  which  have  not  yet 
been  carefully  studied.  Fig.  43  shows  the  form  in  which 
native  sulphur  crystallizes.  Since  native  sulphur  has  re- 
mained as  it  is  for  unlimited  years,  it  follows  that  this 
must  be,  at  ordinary  temperatures,  the  natural  state  of  sul- 
phur, the  stable  form  to  which  all  others  change  if  time 
enough  be  given.  We  shall  call  this  form  a-sulphur. 

That  this  is  not  the  natural  state  of 
sulphur  at  slightly  elevated  temperatures 
we  can  show  by  heating  one  of  the  trans- 
parent crystals  of  a-sulphur  in  a  corked 
test-tube,  dry  inside,  and  surrounded  by 
boiling  water.  It  slowly  becomes  opaque 
and,  while  the  form  remains  unchanged,  the  crystal  is 
now  a  mass  of  prisms  like  that  shown  in  Fig.  44,  but 
very  small.  This  prismatic  sulphur  we  shall  call  ^-sulphur. 
Large  quantities  of  it  can  be  made  by  melting  sulphur  in 
a  crucible,  letting  it  cool  until  a  crust  forms,  piercing  the 
crust,  and  pouring  out  the  liquid  interior.  The  inside  of 


FIG.  44.H3-sulphur. 


THE  OXYGEN  GROUP  237 

the  crucible  is  lined  with  honey-yellow,  transparent  crys- 
tals shaped  like  Fig.  44,  but  elongated  so  as  to  form  needles. 
If  we  allow  the  crucible  to  stand  overnight,  we  find  the  crys- 
tals dull  yellow  and  opaque.  Each  prism  has  changed  to  an 
aggregate  of  crystals  like  that  shown  in  Fig.  43,  but  very 
minute,  and  while  the  prisms  retain  their  form,  they  are 
now  composed  of  a-sulphur,  not  of  y3-sulphur.  In  other 
words,  /3-sulphur  in  the  cold  behaves  exactly  as  a-sulphur 
does  when  heated.  Now  since  a-sulphur  is  the  natural  state 
in  the  cold  and  /3-sulphur  at  100°,  there  must  be  some  tem- 
perature between  at  which  the  change  of  stability  takes 
place.  This  temperature  is  98°,  and  here  a-sulphur,  /3-sul- 
phur,  and  sulphur-vapor  can  exist  in  equilibrium,  just  as 
ice,  liquid  water,  and  water-vapor 
are  in  equilibrium  at  0°. 

There  is  also  an  amorphous 
form  of  sulphur,  insoluble  in  car- 
bon disulphide.  This  separates  as 
a  white  powder  when  a  solution  of 
sulphur  in  carbon  disulphide  is  ex- 
posed to  sunlight  or  a  beam  of 
electric  light. 

When  very  hot  liquid  sulphur  is 

Suddenly     COOled,     Soft     Sulphur     is       FIG.  45.-mparation  of  soft 

obtained.  A  good  method  of  mak- 
ing it  is  shown  in  Fig.  45.  The  sulphur  is  distilled  in  the 
retort,  and  the  hot  liquid  which  condenses  in  the  neck  runs, 
in  a  thin  stream,  into  cold  water.  Soft  sulphur  is  a  trans- 
parent mass  which  has  somewhat  the  consistence  of  rubber, 
but  is  less  tenacious.  Its  color  varies  from  light  yellow  to 
dark  brown.  It  is  unstable,  becoming  brittle  when  kept, 
and  changing  to  a  mass  of  a-sulphur. 

333.  Hydrogen  sulphide,  H2S. — When  hydrogen  is  led 
into  boiling  sulphur  the  two  elements  combine  to  a  slight 
extent,  and  traces  of  hydrogen  sulphide  are  produced.  The 


238  ELEMENTARY  CHEMISTRY 

compound  is  usually  made  by  the  action  of  dilute  hydro- 
chloric or  sulphuric  acid  on  iron  sulphide : 

FeS  +  2HC1  =  FeCl2  +  H2S,  or 
FeS  +  H2S04  =  FeS04  +  H2S. 

Iron  sulphide  is  easily  obtained  by  melting  iron  and  sul- 
phur together.  A  grayish-black,  metallic  mass  is  formed, 
which  is  broken  into  lumps,  placed  in  a  gas-generating  bot- 
tle, and  covered  with  water.  Hydrochloric  or  sulphuric 
acid  is  then  poured  through  the  funnel-tube,  when  hydro- 
gen sulphide  is  at  once  evolved. 

Hydrogen  sulphide  is  a  colorless  gas  with  an  odor  re- 
sembling that  of  rotten  eggs;  in  fact,  the  odor  of  rotten 
eggs  is  due  to  it,  for  during  the  process  of  decay  the  sul- 
phur, which  the  albumin  of  the  egg  always  contains,  is 
liberated  as  H2S.  Hydrogen  sulphide  is  somewhat  soluble 
in  water.  It  has  been  liquefied  and  solidified.  When  ig- 
nited it  burns  with  a  blue  flame,  the  hydrogen  to  water, 
the  sulphur  to  sulphur  dioxide,  and  if  a  cold  plate  is  in- 
troduced into  the  flame,  the  sulphur  no  longer  burns,  but 
is  deposited  as  a  yellow  coating  on  the  plate.  Hydro- 
gen sulphide  is  highly  poisonous.  Air  containing  less 
than  TV  of  1  per  cent  of  it  by  volume  is  fatal  to  small 
animals. 

Many  of  the  sulphides  of  the  metals  are  insoluble  in 
water  and  are  precipitated  when  hydrogen  sulphide  is  passed 
into  a  liquid  in  which  a  salt  of  the  metal  is  dissolved,  thus: 

CuCl2  +  H2S  =  CuS  +  2HC1. 

Many  of  these  sulphides  have  characteristic  colors. 
Zinc  sulphide  is  white  and  antimony  sulphide  orange.  Some 
are  soluble  in  acids  and  are  not  precipitated  until  the  acid 
in  the  solution  is  neutralized.  For  these  reasons  hydrogen 
sulphide  is  much  employed  by  the  chemist  for  detecting  the 
metals  and  separating  them  from  each  other.  On  account 
of  its  unpleasant  odor  and  poisonous  qualities,  various  at- 


THE  OXYGEN  GROUP 


239 


tempts  have  been  made  to  find  a  substitute  for  it,  but  it  still 
remains  a  necessity  in  the  analytical  laboratory. 


334.  Hydrogen  persulphide. 

— Another  compound  of  hydro- 
gen and  sulphur  is  known.  It 
is  a  yellow  oil  which  spontane- 
ously decomposes  into  HaS  and 
sulphur.  Although  it  has  been 
known  for  a  century,  we  are  still 
in  doubt  as  to  its  composition. 
Very  possibly  its  formula  is  HaS2, 
and  it  corresponds  to  hydrogen 
peroxide,  HaOa.  It  is  called 
hydrogen  persulphide. 


FIG.  46.— Preparation  of  sulphur  dioxide 
from  sulphur  and  oxygen. 


335.  Sulphur  dioxide,  S02,  is  the  chief  product  when 
sulphur  burns  in  the  air  or  in  oxygen.  When  sulphur  is 
heated  gently  in  a  bulb-tube,  through  which  a  slow  current 
of  oxygen  from  a  cylinder  is  passed,  it 
burns  with  a  blue,  slightly  luminous 
flame,  and  sulphur  dioxide  can  be  col- 
lected by  downward  displacement  in  a 
cylinder  (Fig.  46). 

Since  the  formula  of  oxygen  is  02, 
it  is  clear  that  when  sulphur  burns  in 
oxygen    no    change    in    volume    takes 
place,  or,  in  other  words,  sulphur  diox- 
ide contains  its  own  volume  of  oxygen. 
For  let  us  suppose  we  have  32  grams 
of    oxygen  =  02  =  22.4    liters.      This 
will  combine  with  32  grams  of  sulphur, 
producing    one    molecular    weight    of 
IGgen  yields  it«  owa  Tot  S02  =  64  grams  =  22.4  liters.     That 
ume  or  sulphur  dioxide,  oxygen  does  produce  an  equal  volume 
of  sulphur  dioxide  when  sulphur  is  burned  in  it  can  be 
shown  by  the   experiment   represented  in   Fig.   47.      The 


240  ELEMENTARY  CHEMISTRY 

vessel  in  the  figure  is  a  small  retort.  At  a  is  a  small  piece 
of  sulphur.  The  liquid  in  the  neck  of  the  retort  and 
in  the  trough  below  is  mercury.  The  retort  is  filled  with 
dry  oxygen.  Then  a  is  heated  gently  with  a  burner  until 
the  sulphur  takes  fire.  At  first  the  level  of  the  mercury 
sinks  because  the  gas  in  the  vessel  is  expanded  by  heat. 
Finally  the  sulphur  flame  is  extinguished  because  the  oxy- 
gen is  all  consumed,  and,  when  the  retort  is  cold,  it  is  found 
that  the  mercury  is  at  just  the  same  level  as  before  the 
combustion.  In  the  language  of  the  atomic  theory,  we  may 
state  this  result  thus:  Every  molecule  of  oxygen  produces 
a  molecule  of  S02,  and  therefore  the  total  number  of  mole- 
cules is  not  altered  by  the  combustion.  Since  equal  volumes 
of  gases  contain  equal  numbers  of  molecules,  the  volume 
also  must  remain  unchanged. 

We  have  seen  that  when  dilute  sulphuric  acid  acts  upon 
zinc,  zinc  sulphate  is  produced  and  hydrogen  liberated. 
Copper  behaves  differently  with  sulphuric  acid.  The  dilute 
acid  does  not  act  upon  it,  but  hot  strong  sulphuric  acid  con- 
verts it  into  copper  sulphate,  liberating  sulphur  dioxide : 
Cu  +  2H2S04  =  CuS04  +  2H20  +  S02. 

The  copper  may  be  heated  with  sulphuric  acid  in  a  flask, 
and  the  sulphur  dioxide  collected  by  downward  displace- 
ment, the  apparatus  being  arranged  in  the  same  way  as  for 
the  production  of  chlorine  (p.  76). 

Sulphur  dioxide  is  a  colorless  gas,  with  the  suffocating 
odor  of  burning  sulphur.  It  is  easily  converted  into  a 
colorless  liquid  by  either  cold  or  pressure,  and  the  liquid 
freezes  on  further  cooling.  Water  dissolves  more  than  thirty 
times  its  volume  of  sulphur  dioxide  at  ordinary  tempera- 
tures. The  solution  is  acid,  and  therefore  must  contain  hy- 
drogen ions.  It  is  believed  that  the  water  and  the  sulphur 
dioxide  unite  to  form  a  compound,  H2S03,  which  then  disso- 
ciates thus :  +  + 

H2SOS  =  H  +  H  +  S03. 


THE  OXYGEN  GROUP 

H2S03  is  therefore  an  acid,  and  the  name  sulphurous 
acid  has  been  given  to  it.  Sulphurous  acid  is  only  known 
in  solution,  but  its  salts,  in  which  the  hydrogen  is  replaced 
by  metals,  are  stable  and  can  readily  be  obtained.  These 
salts  are  called  the  sulphites. 

336.  Sodium  sulphite,  Na2S03,  is  the  most  important. 
It  forms  white  crystals  with  7H20.     It  is  largely  used  in 
photography,  and  to  some  extent  also  as  a  preservative. 
Thus,  when  added  to  fruit  juices,  like  cider,  it  retards  or 
altogether  prevents  the  fermentation  which  would  other- 
wise at  once  take  place.     However,  the  addition  of  any 
chemical  product  to  articles  of  food  is  a  dangerous  opera- 
tion and  ought  to  be  preceded  by  the  most  complete  proof 
of  the  harmlessness  of  the  substance  added.    Except  those 
of  the  sodium  group,  the  sulphites  of  the  metals  are  insolu- 
ble in  water. 

337.  Sulphur  trioxide,    S03. — Sulphur   dioxide   unites 
with  oxygen,  thus, 

S02  +  0  =  S03, 

but  the  reaction  is  extremely  slow.  Certain  substances,  like 
finely  divided  platinum,  or  iron  oxide,  Fe203,  cause  the  re- 
action to  take  place  rapidly.  Sulphur  trioxide  is  made  by 
leading  a  mixture  of  sulphur  dioxide  and  oxygen  over  one 
of  these  substances  at  a  gentle  heat.  In  the  laboratory, 
as  well  as  on  the  large  scale,  asbestos  covered  with  finely 
divided  platinum  is  employed.  This  method  of  making  sul- 
phur trioxide  is  just  now  becoming  of  immense  commercial 
importance.  For  the  S03  produced  when  brought  into  con- 
tact with  water  yields  sulphuric  acid,  H2S04,  and  this  way 
of  making  sulphuric  acid  bids  fair  to  displace  the  older 
method. 

Sulphur  trioxide  appears  to  exist  in  two  modifications. 
One  is  a  transparent  liquid,  which  solidifies  a  little  below 
room  temperature.  When  this  liquid  is  allowed  to  stand 


ELEMENTARY  CHEMISTRY 

it  slowly  passes  into  the  second  modification,  a  mass  of  white 
silky  needles,  resembling  asbestos.  Either  form,  when 
brought  into  contact  with  water,  combines  energetically 
with  it,  producing  a  hissing  noise  and  great  heat : 

S03  +  H20  =  H2S04. 

S03  is  thus  the  anhydride  of  sulphuric  acid,  and  can  be 
called  sulphuric  anhydride. 

338.  Sulphuric  acid,  H2S04. — When  sulphur  dioxide,  air, 
and  water  are  left  in  contact  sulphuric  acid  is  slowly  pro- 
duced— 

S02  +  0  +  H20  =  H2S04— 

but  the  reaction  is  so  slow  that  the  acid  could  not  be  made 
practically  by  this  method.  But  in  presence  of  nitrogen 
peroxide,  N02,  the  reaction  becomes  rapid.  Chemists  are 
still  at  odds  over  the  role  of  the  nitrogen  peroxide.  All 
that  we  can  do  here  is  to  say  that  the  action  is  a  catalytic 
one  (pp.  42-49),  so  that  the  nitrogen  peroxide  is  not  used 
up,  but  is  able  to  produce  large  quantities  of  sulphuric  acid, 
while  remaining  unchanged  itself,  and  this  statement  about 
sums  up  the  present  knowledge  of  the  matter. 

339.  The  lead-chamber  process. — This  accelerating  in- 
fluence of  nitrogen  peroxide  upon  the  combination  of  sul- 
phur dioxide,  oxygen,  and  water  to  sulphuric  acid  is  the 
basis  of  the  lead-chamber  process  for  the  manufacture  of  the 
acid.    The  sulphur  dioxide  is  made  by  burning  pyrite,  FeS2, 
when,  according  to  the  usual  behavior  of  sulphides,  the 
sulphur  burns  to  S02  and  the  iron  remains  as  Fe203.     The 
combustion   is  so   conducted   that   the   S02   contains   free 
oxygen  from  the  air,  which,  as  will  be  seen  at  once  from  the 
equation    (paragraph   338),   is  necessary   for  the  reaction 
which  is  to  follow.     The  hot  gases  from  the  furnace  are 
cooled  and  charged  with  nitrogen  peroxide  by  methods  which 
we  shall  not  discuss.    Then  they  pass  into  a  large  chamber 
built  of  thin  plates  of  lead.    The  size  of  this  chamber  differs 


THE  OXYGEN  GROUP  243 

in  different  works.  It  may  be  100  feet  long,  20  feet  wide, 
and  20  feet  high.  Into  this  chamber  jets  of  steam  from  a 
boiler  are  passed,  and  here  the  combination  of  the  sulphur 
dioxide,  water,  and  oxygen  takes  place  under  the  influence 
of  the  nitrogen  peroxide,  and  the  sulphuric  acid  falls  in  a 
fine  rain  to  the  floor.  Through  a  short  wide  tube  of  lead 
the  gases  pass  into  a  second  similar  chamber,  and  then  into 
a  third,  to  complete  the  reaction. 

Since  more  steam  is  injected  into  the  chambers  than  can 
take  part  in  the  reaction,  the  sulphuric  acid  contains  water, 
usually  about  35  per  cent.  Some  of  this  water  can  be 
evaporated  away  by  heating  the  acid  in  shallow  lead  pans, 
but  when  the  evaporation  has  gone  so  far  that  the  acid 
contains  80  per  cent  H2S04  it  begins  to  attack  the  lead 
and  further  evaporation  must  be  conducted  in  vessels  of 
platinum,  which  are  plated  inside  with  gold,  for  the  strong 
acid  acts  somewhat  even  upon  platinum. 

This  method  of  making  sulphuric  acid  is  carried  out 
on  an  enormous  scale.  About  four  million  tons  of  the  acid 
are  made  by  it  yearly  throughout  the  world,  and  nearly  a 
million  in  the  United  States.  But  just  at  present  it  seems 
to  be  in  serious  danger  of  being  abandoned  in  favor  of  the 
method  spoken  of  just  above  (p.  241),  in  which  sulphur 
trioxide  is  made  first  by  passing  S02  and  oxygen  over  as- 
bestos coated  with  powdered  platinum,  and  then  the  S03 
brought  in  contact  with  water.  Already  some  of  the  largest 
works  in  the  world  have  discontinued  the  lead-chamber  proc- 
ess and  adopted  the  new  method. 

340.  Properties. — Pure  sulphuric  acid  is  a  colorless  oil 
which  freezes  at  10°  to  white  crystals.  When  heated  it 
boils  at  a  high  temperature  (340°),  but  the  suffocating 
vapor  produced  is  not  H2S04.  It  is  a  mixture  of  sulphur 
trioxide  and  steam,  for  sulphuric  acid  decomposes  when  it 
is  vaporized: 

H2S04  =  H20  +  S03. 
17 


244  ELEMENTARY  CHEMISTRY 

This  we  know  from  the  density  of  the  vapor.    The  mo-, 
lecular  weight  of  sulphuric  acid  is  98. 

2TL  =    2 

S  —  32 

40  =  16  X  4  =  64 
"98 

Hence  if  the  vapor  was  H2S04  its  density  referred  to 

98 

hydrogen  would  be  —  =  49. 
Z 

But  if  it  is  separated  into  H20  and  S03  when  volatilized 
the  volume  of  the  vapor  would  be  doubled  and  the  density 
would  be  half  the  above  figure.  When  the  density  of  the 
vapor  is  actually  determined,  it  is  found  to  be  24.5,  and 
this  proves  that  the  splitting  up  referred  to  must  occur. 
We  have  already  studied  a  similar  case  under  ammonium 
chloride  (p.  108). 

Sulphuric  acid  and  water  dissolve  each  other  in  all  pro- 
portions, and  much  heat  is  liberated  when  the  two  liquids 
are  mixed.  The  acid  must  always  be  poured  in  a  thin  stream 
into  the  water,  never  the  reverse.  The  acid  has  a  violent  ac- 
tion upon  organic  tissues,  and  must  be  removed  at  once, 
if  got  on  the  skin,  by  copious  washing  with  water. 

We  have  already  noticed  instances  of  the  action  of  sul- 
phuric acid  upon  the  metals.  Upon  some,  like  iron  and 
zinc,  the  dilute  acid  acts,  liberating  hydrogen.  Upon  others, 
like  copper  and  mercury,  the  dilute  acid  has  little  or  no 
action,  but  the  hot  strong  acid  attacks  the  metal  ener- 
getically, liberating  no  hydrogen,  but  sulphur  dioxide  in- 
stead. 

Most  of  the  sulphates  of  the  metals  are  soluble  in  water. 
We  have  seen  that  barium  sulphate  is  insoluble,  and  that 

its  formation  is  employed  as  a  test  both  for  Ba  ions  and  S04 
ions. 


THE  OXYGEN  GROUP  245 

Many  other  acids  composed  of  hydrogen,  sulphur,  and 
oxygen  are  known.  The  only  one  we  shall  discuss  is 

341.  Thiosulphuric  acid,  H2S203.    This  name  means  sul- 
pho-sulphuric  acid,  and  the  compound  is  so  called  because 
it  may  be  regarded  as  sulphuric  acid  in  which  one  oxygen 
atom  has  been  replaced  by  a  sulphur  atom.    The  acid  itself 
has  never  been  obtained,  but  its  salts  are  well  known.    They 
are  called  the  thio sulphates. 

Sodium  thiosulphate,  Na2S203,  is  the  most  important. 
It  is  produced  when  a  solution  of  sodium  sulphite  is  boiled 
with  sulphur: 

Na2S03  +  S  =  Na2S203. 

It  forms  large  colorless  crystals  containing  five  molecular 
weights  of  water.  It  is  very  soluble  in  water,  and  the  solu- 
tion is  largely  employed  in  the  fixing  process  in  photography 
(p.  157),  and  also  in  the  preparation  of  paper-pulp.  The 
pulp  is  bleached  by  chlorine  and  the  excess  of  chlorine  must 
be  removed,  since  it  would  injuriously  affect  the  paper. 
Sodium  thiosulphate  accomplishes  this.  In  this  way  it 
comes  about  that  most  finished  paper  contains  a  little  so- 
dium thiosulphate. 

342.  Sulphur  chloride,  S2C12,  is  produced  when  chlorine 
is  led  over  melted  sulphur  heated  in  a  retort.    It  is  a  heavy, 
yellow-brown  liquid  with  a  peculiar,  unpleasant  smell.     It 
dissolves  sulphur  freely  and  is  used  in  vulcanizing  rubber, 
in  which  process  the  rubber  takes  up  sulphur.     Ordinary 
soft  rubber  contains  2  to  3  per  cent  of  sulphur ;  hard  rubber 
considerably  more. 

343.  Selenium  and  tellurium,  the  remaining  elements  of  the 
group,  strongly  resemble  sulphur  chemically,  but  are  far  less  abun- 
dant.    Their  hydrogen  compounds  (H2Se  and  HaTe)  are  colorless, 
foul- smelling,  highly  poisonous  gases.     Selenic  acid,  HaSeO4,  corre- 
sponding to  H3SO4,  is  interesting  because  it  is  able  to  dissolve  gold. 


CHAPTER    XXXVI 

MAMGAJfKSS 
Mn  =  55. 

344.  This  metal  bears  about  the  same  relation  to  chlo- 
rine as  chromium  does  to  sulphur.    It  is  a  metal  very  simi- 
lar to  iron,  and  like  chromium  it  enters  into  the  structure 
of  acids,  which,  like  the  corresponding  chromium   com- 
pounds, are  known  only  as  salts. 

345.  Occurrence.— Manganese   carbonate,    MnCO$.    and 
various  oxides  of  the  element  occur  in  nature,  but  the  most 
abundant  and  important  compound  is  manganese  dioxide. 
MnOg.    This  occurs  as  the  mineral  pyrolusite  in  gray-black, 
metallic-looking  crystals  which  yield  a  black  powder.     We 
have  discussed  the  use  of  manganese  dioxide  for  the  prepa- 
ration of  chlorine  (p.  73),  and  it  is  still  much  employed  on 
the  large  scale  for  this  purpose,  but  it  is  almost  certain 
that  the  manufacture  of  chlorine  by  the  electrolysis  of  so- 
dium  chloride   solution   will   shortly  displace  this   older 


Manganese  is  widely  distributed.  Traces  are  contained 
in  sea-water,  in  most  spring  waters,  and  in  many  •"""•1« 
and  plants. 

Preparation. — Minr^::---  i?  :-r?t  or.  tdin-rl  bv  h -rating 
strongly  at  one  point  a  mixture  of  manganese  dioxide  and 
powdered  aluminium.  The  reaction  spreads  through  the 
HJIHP,  which  is  carried  to  an  extremely  high  temperature  by 


MANGANESE  247 

the  heat  liberated  in  the  chemical  change.  Aluminium 
oxide  and  manganese  are  produced: 

3Mn02  +  4A1  =  2A120,  -f  3MiL 

346.  Properties. — Manganese  is  a  reddish-gray  metal 
with  a  strong  luster.  It  is  about  as  hard  as  iron  and,  when 
pure,  is  scarcely  acted  upon  by  the  air.  Even  weak  acids 
readily  dissolve  it,  liberating  hydrogen  and  producing  man- 
ganous salts,  in  which  the  manganese  is  bivalent: 

Mn  +  2H  =  Mn  +  Hj,  or 
Mn  +  2HC1  =  MnCl2  +  H,. 
Manganous 
chloride. 

These  salts  are  the  most  stable  and  important  com- 
pounds in  which  manganese  plays  the  role  of  a  metal.  The 

ion  Mn  communicates  a  delicate  rose-color  to  liquids  con- 
taining it,  and  since  this  color  is  due  solely  to  the  man- 
ganese ion,  the  solutions  of  all  manganous  salts  have  the 
same  color  (p.  131).  Crystallized  manganous  salts  have  the 
same  tint. 

Manganous  sulphate,  MnSO^  is  a  white  powder,  soluble 
in  water.  When  the  solution  is  evaporated,  crystals  sepa- 
rate which  may  contain  1,  4,  5,  or  7  H2O,  according  to  the 
temperature.  As  usual  in  such  cases,  the  higher  the  tem- 
perature of  crystallization  the  less  water  in  the  crystals 
(p.  175).  The  crystals  are  rose-colored,  like  the  other  man- 
ganous salts. 

The  oxide  corresponding  to  these  salts  is,  of  course,  the 
oxide  in  which  manganese  is  bivalent — that  is,  manganous 
oxide,  MnO.  It  is  a  green  powder,  made  by  heating  the  car- 
bonate in  the  absence  of  air : 

MnC03  =  MnO  +  CO2. 

847.  There  is  also  an  oxide  Mn,O,,  in  which  the  metal  is  triva- 
lent.  It  is  a  black  powder.  Corresponding  to  it  is  a  series  of  un- 


248  ELEMENTARY  CHEMISTRY 

stable  salts  which  also  contain  trivalent  manganese — e.  g.,  manganic 

sulphate.   Mn3(SO4)3.     These  are  called  the  manganic  salts.    Their 

+++ 
solutions  contain  the  ion  Mn,  and  are  colored  violet. 

Various  other  oxides  of  manganese  are  known.  We 
shall  only  mention 

348.  Manganese  heptoxide,  Mn207,  which  is  interesting 
because  it  corresponds  to  chlorine  heptoxide,  C1207,  and 
therefore  forms  one  of  the  few  points  of  resemblance  be- 
tween manganese  and  the  halogens.     It  is  a  thick  liquid 
with    a    greenish-black    metallic    luster.      When    carefully 
heated  it  passes  into  a  purple  vapor.    It  is  a  dangerous  sub- 
stance and  often  spontaneously  separates  into  manganese 
dioxide  and  oxygen,  with  energetic  explosion: 

Mn207  =  2Mn02  +  30. 

349.  Permanganic    acid,    HMn04. — When    manganese 
heptoxide  is  dissolved  in  much  cold  water,  the  liquid  ac- 
quires a  beautiful  reddish-violet  color,  and  is  found,  when 
the  usual  tests  are  applied  to  it,  to  contain  an  acid.     A 
chemical    reaction    has    occurred,    and    permanganic   acid, 
HMn04,  is  dissolved  in  the  water: 

Mn207  +  H20  =  2HMn04. 

Permanganic  acid  is  unstable,  and  is  only  known  in 
solution,  but  its  salts  can  readily  be  obtained.  The  most 
important  salt  is 

Potassium  permanganate,  KMn04,  which  forms  almost 
black  crystals,  with  a  greenish  metallic  luster.  It  is  quite 
soluble  in  water,  and  the  solution  has  the  same  magnificent 
red-violet  tint  which  we  have  already  noticed  in  the  water 
solution  of  permanganic  acid,  and  which,  of  course,  is  due 

to  the  ion  Mn04.  It  is  clear  from  the  formula  that  potas- 
sium permanganate  is  rich  in  oxygen.  This  oxygen  is  not 
firmly  held  and  is  easily  given  up  to  any  oxidizable  sub- 


MANGANESE  249 

stance.  Hence  the  solution  of  the  salt  is  a  powerful  oxidiz- 
ing agent  and  is  much  employed  for  that  purpose.  Thus  it 
attacks  and  destroys  most  kinds  of  organic  matter,  produ- 
cing very  much  the  same  effect  in  the  cold  as  free  oxygen  at 
high  temperatures. 


CHAPTER    XXXVII 

THE  HALOGENS 

Fluorine,  F.          Chlorine,  01.          Bromine,  Br.          Iodine,  I. 

350.  These  four  elements,  taken  together,  are  called  the 
halogens.  They  resemble  each  other  strongly.  In  their 
compounds  with  hydrogen  and  the  metals  they  are  univa- 
lent.  The  hydrogen  compounds  (HF,  HC1,  HBr,  and  HI) 
are  colorless,  suffocating  gases,  very  soluble  in  water,  and 
their  solutions  are  strongly  acid.  The  compounds  of  these 
four  elements  with  the  same  metal  usually  crystallize  in 
the  same  form.  Thus  NaF,  NaCI,  NaBr,  and  Nal  all  crys- 
tallize in  cubes.  The  intensity  of  the  color  of  the  elements 
increases  from  fluorine  (F  =  19),  which  is  a  gas  of  a  some- 
what paler  color  than  that  of  chlorine,  down  to  iodine 
(I  =  127),  which  is  almost  opaque,  but  in  a  thin  film  al- 
lows a  little  reddish-brown  light  to  pass.  In  the  same  order 
the  chemical  activity  and  non-metallic  character  decrease. 
Fluorine  is  the  most  energetic  non-metal,  and  by  far 
the  most  active,  chemically,  of  all  the  elements,  and 
the  tendency  of  the  halogens  to  enter  chemical  changes  de- 
creases down  to  iodine,  which,  however,  is  by  no  means 
inert. 

Chlorine,  which  has  already  been  described  (p.  76),  is 
the  most  abundant,  and  fluorine  comes  next.    Bromine  and 
iodine  are  far  less  common.     These  elements  only  occur  in 
compounds  in  nature;  never  native. 
250 


THE  HALOGENS 


251 


FLUORINE,  F-=  19. 

351.  Occurrence. — The  chief  natural  fluorine  compound 
is  fluorspar,  calcium  fluoride,  CaF2,  an  abundant  and  im- 
portant mineral.    It  crystallizes  in  cubes  which  are  colorless 
and  transparent  when  pure,  but  usually  colored  by  impu- 
rities.   Some  samples  of  it  contain,  in  addition  to  the  fluor- 
ine combined  with  the  calcium,  free  fluorine  dissolved  in  the 
crystals  of  calcium  fluoride.     It  escapes  as  gas  when  the 
crystals  are  powdered.     Fluorine,  therefore,  occurs  in  the 
free  state  in  nature,  but  only  in  traces.     Compounds  of 
fluorine  are  widely  distrib- 
uted.    They  are  found  in 

traces  in  river  and  sea 
water,  in  many  plants,  and 
in  the  bones,  teeth,  blood, 
and  brain  of  animals. 

352.  Preparation. — 
Fluorine  is   made   by   the 
electrolysis  of  liquid  anhy- 
drous    hydrofluoric     acid, 
HF.      Since    hydrofluoric 
acid  itself  is  almost  a  non- 
conductor of  the  current, 
some  potassium  hydrogen 


dis- 


FIG.  48.— Isolation  of  fluorine. 


fluoride,  KHF2,  is 
solved  in  it.  It  then  conducts  fairly  well.  This  liquid  is 
placed  in  a  U-shaped  tube  of  copper  (Fig.  48).  This 
tube  is  closed  by  stoppers,  F  F,  made  of  fluorspar,  through 
which  pass  the  platinum  electrodes,  1 1,  which  convey  the  cur- 
rent. Through  side  tubes  the  gaseous  products  of  the  elec- 
trolysis escape  and  can  be  examined.  The  copper  tube  is  im- 
mersed in  a  freezing  mixture,  for  the  electrolysis  must  be 
carried  out  at  a  low  temperature  to  prevent  the  vaporization 
of  the  hydrofluoric  acid.  Equal  volumes  of  hydrogen  and 


252  ELEMENTARY  CHEMISTRY 

fluorine  are  produced.  The  hydrogen  is  liberated  from  the 
negative  pole  and  the  fluorine  at  the  positive. 

353.  Properties. — Fluorine  is  a  gas  of  a  color  somewhat 
similar  to  that  of  chlorine,  but  paler.  Its  odor  is  distinct 
from  that  of  chlorine,  but  it  has  the  same  irritating  action 
upon  the  mucous  membranes.  When  the  gas  is  passed  into 
a  vessel  cooled  by  being  surrounded  by  liquid  air  it  condenses 
to  a  yellow  liquid  which  is  slightly  heavier  than  water. 

The  chemical  activity  of  fluorine  is  extraordinary,  and 
in  this  respect  it  easily  stands  first  among  the  elements. 
All  of  the  metals — even  gold  and  platinum — are  attacked  by 
it  and  converted  into  fluorides.  In  the  case  of  copper  the 
action  is  slight,  and  for  this  reason  copper  is  employed  in 
the  construction  of  the  apparatus  used  in  the  preparation 
of  the  element.  Sulphur,  phosphorus,  and  most  of  the  other 
non-metals  ignite  spontaneously  in  the  gas,  and  burn  to  the 
corresponding  fluorides.  Even  bromine  and  iodine,  which 
are  very  similar  to  fluorine  chemically,  take  fire  in  it  and 
burn  energetically,  and  this  fact  is  worth  noting,  for,  as  a 
rule,  elements  which  closely  resemble  each  other  in  their 
chemical  conduct  show  little  tendency  to  unite.  Fluorine 
appears  to  have  no  action  upon  chlorine  nor  upon  oxygen. 
When  a  tube  from  which  fluorine  is  escaping  is  brought 
into  an  atmosphere  of  hydrogen  there  is  an  explosion,  and 
the  fluorine  burns  to  hydrofluoric  acid,  HF,  producing  a  hot 
blue  flame.  This  occurs  in  the  dark  as  well  as  in  daylight. 
(Compare  the  behavior  of  chlorine  and  hydrogen,  p.  78.) 

Fluorine  also  acts  violently  upon  many  compounds. 
When  it  is  passed  into  a  test-tube  containing  a  little  dry 
salt  (NaCl)  at  the  bottom,  the  tube  at  once  becomes  full 
of  chlorine,  sodium  fluoride  being  produced: 

NaCl  +  F  =  NaF  +  Cl. 

In  the  same  way  it  acts  upon  other  chlorides.  Its  action 
upon  iodides  is  similar,  but  the  iodine,  which  is  liberated 


THE   HALOGENS  253 

at  first,,  burns  afterward  to  iodine  fluoride.  It  reacts  with 
water  instantly,  even  in  the  dark  (compare  the  same  reac- 
tion with  chlorine,  p.  79),  forming  hydrofluoric  acid  and 
liberating  oxygen  as  ozone.  Its  action  upon  other  hydrogen 
compounds  is  similar.  Thus,  when  liquid  fluorine  is  spilled 
upon  the  floor,  a  flame  rises  which  is  due  to  the  combination 
of  the  fluorine  with  the  hydrogen  of  the  wood.  Fluorine 
does  not  act  upon  glass,  and,  if  it  is  perfectly  pure,  it  can 
be  preserved  sealed  up  in  glass  tubes. 

354.  Hydrofluoric  acid,  HF,  is  produced  with  explosion, 
when  hydrogen  and  fluorine  are  brought  together.  It  is 
made  by  the  action  of  strong  sulphuric  acid  on  powdered 
calcium  fluoride  at  a  gentle  heat: 

CaF2  +  H2S04  =  CaS04  +  2HF. 

Since  the  acid  acts  rapidly  upon  glass,  this  operation 
must  be  conducted  in  vessels  of  platinum  or  of  lead.  If  a 
water  solution  of  the  acid  is  required,  the  mixture  of  cal- 
cium fluoride  and  sulphuric  acid  is  heated  in  a  lead  retort, 
and  the  gaseous  HF  brought  into  contact  with  water  con- 
tained in  a  platinum  dish.  If,  on  the  other  hand,  it  is  de- 
sired to  obtain  the  acid  free  from  water,  the  gas  is  led  into 
a  platinum  bottle  surrounded  by  a  freezing  mixture.  Hy- 
drofluoric acid  is  thus  obtained  as  a  colorless  liquid,  which 
must  be  kept  in  a  freezing  mixture,  since  otherwise  it  will 
vaporize  and  burst  the  bottle,  if  the  latter  is  sealed. 

Hydrofluoric  acid  is  a  colorless,  poisonous  gas,  with  a 
suffocating  odor.  It  is  freely  soluble  in  water  to  a  colorless 
fuming  liquid,  which  is  strongly  acid  and  readily  dissolves 
zinc,  iron,  silver,  and  many  other  metals,  liberating  hydro- 
gen and  forming  the  corresponding  fluorides.  Commercial 
hydrofluoric  acid  is  the  water  solution  of  the  gas.  On  ac- 
count of  its  action  upon  glass  it  is  sold  in  bottles  of  hard 
paraffin,  upon  which  it  has  no  action. 

Over  a  little  finely  powdered  silica,  Si02,  contained  in  a 


254 


ELEMENTARY  CHEMISTRY 


platinum  crucible  we  pour  carefully  some  strong  hydro- 
fluoric acid.  There  is  a  violent  reaction,  heat  is  produced, 
the  white  powder  disappears,  and  a  colorless  suffocating 
gas  escapes,  which  produces  fumes  in  the  air.  This  gas  is 
silicon  fluoride,  SiF4,  produced  thus : 

Si02  +  4HF  =  SiF4  +  2H20. 

This  experiment  will  help  us  to  understand  the  action 
of  hydrofluoric  acid  on  glass.1  For  etching  on  glass,  gaseous 
HF  is  usually  employed.  In  marking  the  graduations  on  a 
thermometer  tube,  for  instance,  the  tube  is  first  coated  thin- 
ly with  wax  and  then  the  divisions  scratched  through  the 
coating,  laying  bare  the  glass.  In  this  condition  the  tube 
is  exposed  for  a  time  to  hydrofluoric  acid  gas,  produced  by 
a  mixture  of  calcium  fluoride  and  strong  sulphuric  acid. 
When  the  wax  is  removed  the  divisions  are  found  etched  on 
the  glass. 

The  water  solution  of  hydrofluoric  acid  acts  violently 
upon  the  skin,  and  care  must  be  taken  not  to  get  it  upon 
the  hands.  Its  vapor  is  poisonous  and  must  not  be  inhaled. 

Chlorine  has  already  been  described  (Chapter  X). 


,«. 

II 

tf.S 

f-o, 


a  ain  th^  paragraphs  on  the  composition  of  glass  (p.  199). 

r 9  \AxM^l  OO^O^CA/! 
6 


CHAPTER    XXXVIII 

BROMINE   AND    IODINE 

BROMINE,  Br  =  80. 

355.  Some   potassium   bromide,   KBr,   is  dissolved  in 
water  and  a  little  chlorine  gas  passed  into  the  liquid.    At 
once  the  red  color  of  bromine  appears  in  the  solution  and, 
if  much  potassium  bromide  was  used  and  the  current  of 
chlorine  be  continued,  drops  of  bromine  will  collect  under 
the  solution.     The  solution  of  potassium  bromide  contains 

bromine  ions  Br,  and  the  tendency  of  chlorine  to  exist  as  an 
ion  is  greater  than  that  of  bromine,  so  that  the  reaction 

2Br  +  C12  =  2C1  +  Br2 

occurs,  and  bromine  separates. 

356.  Occurrence  and  preparation. — The  occurrence  of 
great  deposits  containing  potassium  chloride  at  Stassfurt, 
in  Germany,  has  been  mentioned.    This  potassium  chloride 
is  separated  from  other  substances  which  occur  along  with 
it  by  treating  the  deposit  with  water  and  then  crystallizing 
the  potassium  chloride  from  the  solution.    Now  the  liquid 
from  which  the  potassium  chloride  has  separated  contains 
bromine  ions,  and  when  treated  with  chlorine,  as  in  the 
experiment  just  described,  bromine  is  obtained.    The  liquid 
is  hot  when  the  chlorine  is  introduced,  and  the  bromine 
escapes  as  a  vapor,  which  is  condensed.     About  two-thirds 
of  the  bromine  of  the  world  is  made  at  Stassfurt.     The 
rest  is  extracted  by  similar  methods  from  the  waters  of 

255 


256  ELEMENTARY  CHEMISTRY 

salt  springs  in  the  United  States.  The  water  of  the  Dead 
Sea  contains  bromine  ions,  and  is  likely  to  be  employed  for 
the  production  of  the  element  in  the  near  future. 

357.  Properties. — Bromine  is  a  black-red,  almost  opaque 
liquid,  about  three  times  as  dense  as  water.  When  slightly 
cooled  it  solidifies  to  a  mass  somewhat  resembling  iodine. 
At  ordinary  temperatures,  and  more  rapidly  when  heated, 
the  liquid  gives  off  red  bromine  vapor,  the  odor  of  which  re- 
calls that  of  chlorine.  Its  action  upon  the  eyes,  nose,  and 
throat  is  even  more  violent  than  that  of  chlorine.  The 
density  of  bromine  vapor  is  about  SO  referred  to  hydrogen, 
and  this  shows  that  the  molecular  weight  is  about  160. 
Hence  the  formula  is  Br2 — the  bromine  molecule  contains 
two  atoms.  This  is  also  the  case  with  the  other  halogens. 

Bromine  is  an  energetic  element.  It  unites  directly 
with  nearly  all  of  the  metals,  forming  the  corresponding 
bromides.  Sometimes  the  combination  is  violent.  A  small 
fragment  of  potassium  dropped  into  bromine  produces  a 
loud  explosion.  Curiously  enough,  it  is  quite  without  action 
on  sodium,  which  is  so  similar  to  potassium.  When  a  hy- 
drogen flame  is  lowered  into  bromine  vapor  it  continues  to 
burn,  producing  colorless,  suffocating  hydrobromic  acid, 
HBr.  Bromine  and  oxygen  have  no  action  upon  each  other, 
and  no  compound  of  the  two  elements  has  ever  been  ob- 
tained. Bromine  is  soluble  in  water,  and  the  red  solution 
produced  is  called  bromine  water  and  is  used  in  the  labora- 
tory. Of  course  this  solution  does  not  contain  bromine  ions. 
It  contains  bromine  as  Br2.  Bromine  ions  can  only  exist 

in  presence  of  an  equivalent  number  of  positive  ions,  like 

+  + 

those  of  K  or  Na.  When  bromine  water  is  placed  in  sun- 
light, the  following  reaction  slowly  takes  place : 

Br2  +  H20  =  2HBr  +  0, 

and  oxygen  collects  in  the  vessel. 

Bromine  is  a  splendid  disinfectant,  and  is  employed  for 


BROMINE  AND  IODINE  257 

this  purpose,  but  not  largely,  because  it  is  too  expensive. 
The  most  important  compound  of  bromine  is 

358.  Potassium  bromide,  KBr,  which  crystallizes  in  col- 
orless cubes  which  have   a  pleasant  saline  taste  and  are 
freely  soluble  in  water.    It  is  largely  employed  in  medicine 
in  diseases  of  the  nervous  system,  and  also  in  photography. 
When  added  to  the  developer  (p.  156)  it  prevents  the  produc- 
tion of  an  image  upon  those  parts  of  the  plate  which  have 
only  slightly  been  affected  by  light.    It  is  employed  there- 
fore in  developing  plates  which  have  accidentally  been  ex- 
posed too  long  in  the  camera. 

359.  Hydrobromic   acid,    HBr,   corresponds    to    hydro- 
chloric acid.    It  results  when  dilute  sulphuric  acid  acts  upon 
the  bromide  of  a  metal — on  KBr,  for  instance.1    The  equa- 
tion is 

2KBr  +  H2S04  =  K2S04  +  2HBr. 

Another  method  of  making  hydrobromic  acid  is  to  pass 
a  mixture  of  hydrogen  and  bromine  vapor  over  red-hot 
platinum.  This  mixture  is  easily  obtained  by  passing  hy- 
drogen through  liquid  bromine.  The  bottle  containing  the 
bromine  is  placed 
in  warm  water, 
which  causes  the 
bromine  to  va- 
porize abundant- 
ly. The  hydro- 
gen charged  with  F"  ^™*  br°mine  * 
bromine  vapor 

passes  into  a  wide  glass  tube,  which  contains  a  spiral  of 
platinum  wire  kept  at  a  red  heat  by  the  passage  of  an  elec- 
tric current.  Here  the  combination  takes  place,  and  the 

1  This  method  of  making  HBr  is  excellent  and  has  the  advantage 
of  perfect  analogy  to  the  preparation  of  hydrochloric  acid  by  the 
action  of  sulphuric  acid  on  salt.  The  usual  statement  that  it  is  impos- 
sible to  obtain  HBr  by  a  similar  reaction  is  quite  without  foundation. 


258 


ELEMENTARY  CHEMISTRY 


hydrobromic  acid  can  be  passed  into  water  and  a  solution 
of  it  obtained.1  The  arrangement  of  the  apparatus  is  shown 
in  Fig.  49. 

Hydrobromic  acid  is  similar  to  hydrochloric  acid.  It  is 
a  colorless,  suffocating  gas,  which  produces  fumes  in  the 
air  and  has  a  pungent,  irritating  odor.  It  is  excessively 
soluble  in  water,  and  the  solution  is  strongly  acid,  for  a  very 
large  proportion  of  the  dissolved  HBr  is  dissociated  into 

H  and  Br  ions. 

IODINE,  I  =  127. 

360.  Occurrence  and  preparation.  —  Traces  of  iodine 
compounds  are  contained  in  sea-water,  and  larger  quantities 

in  sea-weeds,  espe- 
cially in  those 
which  grow  at 
great  depths.  The 
manufacture  of  io- 
dine from  sea- 
weeds is  carried  on 
in  Scotland  and 
France.  The  weeds 
are  first  burned, 
and  by  systematic- 
ally treating  the 
ashes  with  water  a 
liquid  is  finally  ob- 

tained which  contains  the  iodides.  This  is  mixed  with 
manganese  dioxide  and  sulphuric  acid  and  distilled.  The 
vapor  of  iodine  is  liberated  and  is  condensed  in  pear-shaped 
clay  receivers  : 

Mn02  +  3H2S04  = 

2H20  +  I2. 


FIG.  50.— Extraction  of  iodine. 


2NaHS04  +  MnS04 


1  This  experiment  is  due  to  Newth,  Chemical  Lecture  Experiments, 
p.  104 


BROMINE  AND  IODINE  259 

The  apparatus  used  is  shown  in  Fig.  50.  The  iodine  is 
purified  by  distillation. 

Small  quantities  of  potassium  iodate,  KI03,  correspond- 
ing to  potassium  chlorate,  KC103,  are  contained  in  the 
sodium  nitrate  deposits  of  Chili  (p.  116).  The  preparation 
of  iodine  from  this  source  has  become  of  great  importance 
and  threatens  soon  to  displace  altogether  the  production  of 
iodine  from  sea-weed. 

361.  Properties. — Iodine   crystallizes    in    shining   gray- 
black  plates,  which  have  the  color  and  luster  of  graphite. 
It  melts  to  a  thick  black  liquid  slightly  above  the  boiling- 
point  of  water  and  boils  at  a  higher  temperature.     When 
mixed  with  air  the  vapor  of  iodine  has  a  magnificent  violet 
color,  but-the  pure  vapor  is  deep  blue.   Its  density  shows  that 
the  formula  is  I2.    It  is  probable  that  all  substances  which 
vaporize  at  all,  do  so  at  all  temperatures,  but  at  high  tem- 
peratures the  vaporization  is  rapid,  and  therefore  noticeable. 
Iodine  is  a  good  example  of  this.    If  a  little  iodine  is  placed 
in  a  bottle  and  the  stopper  inserted,  the  space  above  the 
solid  will  soon  be  colored  distinctly  violet  by  the  vapor, 
and,  after  a  time,  small  glittering  crystals  of  iodine  will 
condense  upon  the  glass  in  the  upper  portion. 

Iodine  is  only  very  slightly  soluble  in  water,  but  it  dis- 
solves freely  in  alcohol,  ether,  chloroform,  and  carbon  disul- 
phide.  The  alcoholic  and  ethereal  solutions  are  reddish- 
brown,  while  the  solutions  in  chloroform  and  carbon  disul- 
phide  have  a  violet  tint  resembling  that  of  iodine  vapor. 

Iodine  is  less  active  chemically  than  the  other  halogens, 
but  still  quite  energetic.  It  unites  with  most  of  the  metals, 
especially  when  heated,  and  in  some  cases — with  potassium, 
for  instance — the  combination  takes  place  with  violent  en- 
ergy. With  hydrogen  it  only  unites  at  a  high  temperature, 
and  then  only  partially. 

362.  Hydriodic  acid,  HI,  is  a  colorless,  fuming,  suffo- 
cating gas  which  dissolves  very  freely  in  water  to  a  strongly 

18 


260  ELEMENTARY  CHEMISTRY 

acid  liquid.     It  resembles  HC1  closely,  but  is  much  more 
readily  decomposed  by  heat. 

363.  Potassium  iodide,  KI,  is  the  most  important  salt 
of  hydriodic  acid,  and,  in  fact,  the  most  important  iodine 
compound  technically.     It  forms  colorless  cubical  crystals 
which  dissolve  readily  in  water.     It  is  much  employed  in 
medicine. 

364.  Only  one  compound  of  iodine  and  oxygen  is  known  with 
certainty.     This  is 

Iodine  pentoxide,  IaO6,  a  white  powder  which  is  made  by  gently 
heating  iodic  add,  HIO3  (analogous  to  HC1OS).  It  dissolves  in  water, 
and  the  solution  is  found  to  contain  iodic  acid — 

Ia06  +  H2O  =  2HIO3. 
Iodine  pentoxide  is  easily  separated  into  iodine  and  oxygen  by  heat. 

Many  salts  are  known  in  which  the  hydrogen  of  iodic  acid  is 
replaced  by  metals.  These  are  called  iodates.  Potassium  iodate, 
KIO8,  has  been  mentioned  (p.  259). 

365.  Free  iodine  produces  an  intense  blue  color1  with 
starch  paste.    The  nature  of  the  blue  substance  still  remains 
unknown.     The  test  is  extremely  delicate.     We  have  seen 
that  the  mixture  of  starch  paste  with  an  iodide,  like  KI, 
can  be  used  as  a  test  for  ozone,  or,  in  fact,  for  anything  which 
will  set  free  the  iodine  (p.  45). 

366.  General  remarks. — The  halogens  form  one  of  the 
best  examples  of  a  natural  family  or  group  of  elements. 
The  sodium  group  (p.  135)  is  another  remarkable  instance 
of  the  natural  relationship  of  a  number  of  elements.     The 
student  should  compare  these  two  groups  with  each  other. 
He  will  find  that  in  each  group  the  resemblance  of  any  ele- 
ment to  the  other  elements  of  the  same  group  is  striking. 
On  the  other  hand,  the  comparison  of  sodium  with  fluorine, 
or  of  potassium  with  chlorine,  will  give  him  a  vivid  and  ac- 
curate idea  of  the  nature  of  the  differences  between  metals 
and  non-metals — that  is,  between  positive  and  negative  ele- 
ments. 


CHAPTER   XXXIX 

IRON 

Fe  =  56. 


1367.  Iron  is  very  similar  to  manganese,  which  precedes 
t  in  the  order  of  increasing  atomic  weights  (Mn  =  55). 
On  the  other  hand,  it  has  many  points  of  resemblance  with 
nickel  and  cobalt,  which  follow  it  (Ni  =  58.7,  Co  =  59). 
Iron  and  nickel  occur  together  in  meteorites,  and  cobalt  and 
nickel  are  always  associated  in  nature.  All  three  metals 
are  strongly  attracted  by  the  magnet.  Iron  is  by  far  the 
most  abundant  and  important. 

368.  Occurrence. — Iron  was  employed  for  weapons  and 
other  cutting  implements  by  prehistoric  man.  The  native 
metal  is  rare  (except  in  meteorites),  but  iron  compounds  are 
so  common  that  the  metal  stands  fourth  in  order  of  abun- 
dance. The  oxides  and  hydroxides  of  iron  are  its  most 
important  ores.  Iron  disulphide,  FeS2,  is  the  very  common 
mineral  pyrite.  Various  other  iron  compounds  enter  large- 
ly into  the  structure  of  rocks.  It  is  safe  to  assume  that 
whenever  a  rock  or  a  soil  is  colored  red,  yellow,  or  green,  the 
color  is  due  to  iron  in  some  form.  Rocks  are  often  colored 
red  by  ferric  oxide,  Fe203,  yellow  by  ferric  hydroxide, 
Fe(OH)3,  and  green  by  various  silicates  of  iron.  Iron  is 
only  contained  in  small  quantities  in  the  bodies  of  animals 
and  plants,  but  nevertheless  it  plays  an  important  role.  In 
both  it  is  connected  with  the  process  of  respiration.  The 
green  parts  of  plants,  in  which  the  carbon  dioxide  of  the  air 
is  decomposed,  the  carbon  built  into  the  structure  of  the 

OA1 


262  ELEMENTARY  CHEMISTRY 

plant  and  the  oxygen  returned  to  the  atmosphere,  contain 
iron,  and  it  has  been  shown  that  if  no  iron  compounds  are 
supplied  to  the  plant  the  green  substance  is  no  longer  formed 
and  respiration  ceases.  In  animals  the  iron  is  contained  in 
the  red  blood-corpuscles,  by  means  of  which  oxygen  is  taken 
from  the  air. 

369.  Preparation  and  properties. — Pure  iron  can  be  ob- 
tained by  heating  pure  ferric  oxide  in  hydrogen: 

Fe203  +  3H2  =  3H20  +  2Fe. 

Pure  iron,  in  a  compact  mass,  has  a  silver-white  metallic 
luster  which  is  permanent  in  dry  air,  but  rapidly  lost  under 
the  action  of  air  and  water  together.  At  the  same  time  a 
yellow-brown  coating  called  "  rust "  appears  on  the  iron. 
This  contains  ferric  hydroxide  and  ferric  oxide,  and  has  the 
following  composition: 

Fe2032Fe(OH)3. 

When  rusting  has  once  begun  it  proceeds  more  and  more 
rapidly  until  the  change  is  complete. 

When  pure,  iron  melts  only  at  1800°,  which  is  very  much 
above  a  white  heat  (about  1200°).  At  the  temperature  of 
the  electric  arc  it  boils,  rapidly  disappearing  as  vapor. 
When  heated  to  redness  in  the  air,  the  surface  of  the  metal 
is  converted  into  an  oxide,  and  iron  will  burn  brilliantly  in 
oxygen  at  high  temperatures.  Very  finely  powdered  iron 
takes  fire  spontaneously  in  the  air.  Iron  dissolves  easily 
in  acids. 

370.  Preparation  of  iron  on  the  large  scale.— This  is  the 
object  of  the  BLAST-FURNACE  PROCESS.     The  principle  of 
the  process  is  simple  and  is  one  that  we  have  repeatedly 
met.     At  high  temperatures  carbon  acts  on  iron  oxide  as 
it  does  on  other  metallic  oxides,  converting  it  into  metal. 
The  blast  furnace  may  be  100  feet  high,  and  its  interior  is 
egg-shaped,  so  that  the  widest  part — which  may  be  25  feet 


ROBERT  WILHELM   BUNSEN 
B.  Germany,  1811.     D.  1899. 


IRON  263 

across — is  much  nearer  the  bottom  than  the  top.  At  the 
top  the  ore  is  introduced,  together  with  some  form  of  car- 
bon, often  coke.  At  the  bottom  the  liquid  iron  collects  and 
is  allowed  to  flow  out  at  intervals  and  received  in  a  bed  of 
sand,  which  is  arranged  beforehand  in  such  a  way  that  the 
iron  flows  into  small  hollows  provided  for  it,  and  solidifies 
there  to  pig-iron.  A  blast  of  hot  air  under  some  pressure 
is  introduced  near  the  bottom  of  the  furnace,  and  the  burn- 
ing of  the  carbon  by  the  oxygen  supplied  in  this  way  keeps 
the  temperature  of  this  portion  at  a  bright-yellow  heat. 
The  upper  portions  are  not  so  hot. 

This  is  an  outline  of  the  process  as  it  might  be  worked 
with  pure  iron  oxide  to  produce  a  pure  iron.  Now  come 
some  complications.  No  actual  iron  ore  is  pure  Fe203.  All 
ores  contain  a  variety  of  impurities,  of  which  the  most 
important  to  us  at  present  are  silica,  Si02,  and  aluminium 
oxide,  A1203.  Along  with  the  coke  and  ore,  limestone  is 
added  at  the  top  of  the  furnace.  The  lime  of  the  limestone 
and  the  silica  and  A1203  of  the  ore  unite  to  produce  a  sili- 
cate of  calcium  and  aluminium,  which  melts  readily  and  col- 
lects as  a  liquid  mass  over  the  melted  iron.  This  is  called 
the  slag.  It  is  allowed  to  run  out  by  an  orifice  called  the 
slag-hole,  which  is  at  a  higher  level  than  the  hole  through 
which  the  melted  iron  is  tapped.  The  iron  obtained  in  this 
way  is  never  pure.  It  contains  from  2  per  cent  upward  of 
carbon,  some  silicon,  and  smaller  quantities  of  phosphorus, 
sulphur,  and  other  substances.  It  is  called  pig-iron  or  cast 
iron. 

Wrought  iron,  or  malleable  iron,  is  far  purer  than  pig- 
iron.  It  is  made  from  pig-iron  by  the  puddling  process. 
This  consists  in  heating  the  pig-iron  to  a  high  temperature 
in  contact  with  ferric  oxide,  Fe203,  and  in  an  atmosphere 
containing  abundant  oxygen  from  the  air.  Partly  by  the 
oxygen  of  the  ferric  oxide  and  partly  by  that  of  the  sur- 
rounding atmosphere  the  carbon  is  burned  to  CO,  carbon 


264  ELEMENTARY  CHEMISTRY 

monoxide,  which  escapes  as  gas,  and  the  silicon  to  Si02, 
which  goes  into  the  slag.  At  the  same  time  the  phos- 
phorus, sulphur,  and  other  impurities  are  also  removed 
by  oxidation.  Wrought  iron  is  softer  and  more  tenacious 
than  cast  iron,  and  melts  at  a  higher  temperature.  Before 
it  melts  it  softens  and  becomes  somewhat  plastic,  and  this 
is  a  most  important  property,  for  it  makes  it  possible  to 
hammer  a  piece  of  iron  into  any  desired  shape  or  to  weld 
two  pieces  together. 

371.  Steel  contains  more  carbon  than  wrought  iron,  but 
less  than  cast  iron.    The  finest  steel — called  crucible  steel — 
is  made  by  heating  bars  of  wrought  iron  for  a  week  or  more 
in  contact  with  powdered  charcoal.     The  iron  slowly  takes 
up  carbon,  and  each  bar — even  in  the  interior — is  converted 
into  steel.     The  bars  are  afterward  melted  in  crucibles  in 
order  to  secure  uniformity  in  the  product.     Such  steel  is 
used  for  the  finer  grades  of  cutlery. 

Steel  is  harder  and  even  more  tenacious  than  wrought 
iron,  and  melts  at  a  lower  temperature.  Another  important 
difference  is  the  fact  that  steel  can  be  tempered,  which  is 
not  the  case  with  iron,  either  wrought  or  cast.  The  cheaper 
varieties  of  steel — rail  steel,  bridge  steel,  and  .so  on — are 
made  in  enormous  quantities  by  other  processes  which  can 
not  be  discussed  here.  They  are  described  completely  in 
Works  on  metallurgy.  Some  idea  of  the  importance  of  iron 
and  steel  in  modern  life  may  be  gained  from  the  fact  that  in 
1903  nearly  18,000,000  tons  of  pig-iron  were  produced  in 
the  United  States  alone.  This  corresponds  to  over  500 
pounds  for  every  inhabitant  of  the  country. 

372.  Ferrous  oxide,  FeO,  is  a  black  powder  which  easily 
takes  up  oxygen  in  the  air,  passing  into  ferric  oxide.    Fer- 
rous hydroxide,  Fe(OH)2,  is  white,  and,  when  exposed  to 
air,  turns  first  green  and  then  red-brown,  being  converted 
into  ferric  hydroxide,  Fe(OH)3.     Both  ferrous  oxide  and 
ferrous   hydroxide   dissolve  in   acids   yielding   the  ferrous 


IRON  265 

salts  in  which  iron  is  bivalent.  The  formulas  of  the  ferrous 
correspond  therefore  to  those  of  the  corresponding  com- 
pounds of  magnesium  or  zinc.  For  example : 

Ferrous  chloride,  FeCl2. 
Ferrous  bromide,  FeBr2. 
Ferrous  nitrate,  Fe(N03)2. 
Ferrous  sulphate,  FeS04. 

Ferrous  salts  are  produced  when  iron  dissolves  in  acids ; 
thus: 

Fe  +  H2S04  =  FeS04  +  H2. 

They  are  green  when  crystallized  and  pale  green  or  col- 
orless when  dissolved  in  water.  The  most  important  fer- 
rous salt  is 

373.  Ferrous  sulphate,  FeS04,  which  can  be  obtained,  as 
indicated  above,  by  the  action  of  dilute  sulphuric  acid  on 
iron.    It  can  be  more  cheaply  made  at  present  by  the  spon- 
taneous oxidation  of  ferrous  sulphide,  FeS,  under  the  influ- 
ence of  air  and  water.    It  forms  green  crystals  with  7H20, 
which  have  long  been  known  as  "  green  vitriol "  or  "  copper- 
as."   It  is  an  excellent  disinfectant. 

All  the  ferrous  salts  absorb  oxygen  and  pass  into  ferric 
salts  when  preserved.  For  this  reason  they  must  be  kept  in 
well-closed  bottles. 

374.  Ferrous  sulphide,  FeS,  is  produced  as  a  black  pre- 

++ 
cipitate  when  ferrous  ions,  Fe,  and  sulphur  ions,  S,  are 

brought  into  the  same  solution,  i.  e.,  when  a  solution  of  a 
ferrous  salt  is  mixed  with  one  of  a  soluble  sulphide — 

FeCl2  +  Na2S  =  2NaCl  +  FeS. 

The  iridescent  film  which  forms  over  the  surface  of 
stagnant  water  in  pools  and  ditches  consists  of  ferrous  sul- 
phide. It  is  readily  obtained  as  a  grayish-black,  dense,  metal- 
lic, fused  mass  by  melting  iron  and  sulphur  together.  It  is 


266  ELEMENTARY  CHEMISTRY 

employed  in  the  laboratory  in  the  production  of  hydrogen 
sulphide  (p.  237). 

375.  Ferric  oxide,  Fe203,  occurs  in  nature  as  the  mineral 
hcematite  in  steel-gray  crystals,  which  leave  a  red  streak 
when  drawn  over  an  unglazed  porcelain  plate,  and  yield  a 
red  powder  when  pulverized  in  a  mortar. 

Ferric  hydroxide,  Fe(OH)3,  is  precipitated  in  rust-col- 
ored flakes  when  a  solution  of  a  hydroxide  is  added  to  one 
containing  a  ferric  salt — 

FeCl3  +  3NaOH  =  Fe(OH)3  +  3NaCl. 
Ferric  chloride. 

Colloidal  solutions. — Besides  the  ordinary  insoluble  ferric  hydrox- 
ide there  is  another  modification  of  it  which,  when  treated  with 
water,  apparently  dissolves,  forming  a  clear  dark-red  liquid,  which 
leaves  no  residue  when  passed  through  a  filter.  We  have  seen  that 
ordinary  solutions — like  that  of  common  salt  in  water  boil  at  a 
higher  and  freeze  at  a  lower  temperature  than  water  alone.  But 
this  solution  of  ferric  hydroxide  boils  at  100°  and  freezes  at  0°,  just 
as  though  the  water  was  free  from  dissolved  substance.  This  looks 
as  though  the  ferric  hydroxide  was  not  really  dissolved  in  the  water, 
but  only  suspended  in  particles  too  fine  to  be  detained  by  a  filter, 
and  this  idea  is  strongly  supported  by  the  fact  that  when  a  beam  of 
electric  light  is  passed  through  the  liquid  in  a  dark  room,  the  track 
of  the  beam  is  distinctly  visible,  which  is  not  the  case  with  true 
solutions.  This  indicates  the  presence  of  minute  solid  particles 
which  reflect  the  light  to  the  eye  and  make  it  visible,  just  as  the 
path  of  a  sunbeam  is  visible  in  dusty  air.  Many  other  substances 
ordinarily  insoluble  in  water  have  been  obtained  in  a  similar  condi- 
tion. Among  these  are  sulphur  and  many  of  the  metals — for  instance, 
platinum,  silver,  and  gold.  This  curious  state  of  things  is  called 
colloidal  solution,  and,  for  the  reasons  just  stated,  it  is  regarded  at 
present  as  nothing  but  suspension,  in  which  the  solid  particles  are 
very  small.  The  colloidal  solutions  of  the  metals  are  clear,  deep 
brownish  liquids  from  which,  when  they  are  preserved,  the  metal 
gradually  settles  as  a  powder. 

376.  Ferric  salts. — Ferric  oxide  dissolves  with  difficulty 
and  ferric  hydroxide  with  ease  in  acids,  both  producing 


IRON  267 

ferric  salts,  in  which  the  iron  is  trivalent.  Ferric  salts 
which  contain  water  of  crystallization  are  brownish-yellow 
in  the  solid  state,  and  their  solutions  have  the  same  color. 

Ferric  chloride,,  FeCl3,  has  been  found  in  the  craters  of 
volcanoes.  It  is  made  by  burning  piano-wire  (the  purest 
commercial  form  of  iron)  in  a  rapid  current  of  dry  chlorine. 
It  forms  iron-black,  glittering,  tabular  crystals,  which  are 
very  soluble  in  water.  When  the  solution  is  evaporated,  and 
then  allowed  to  crystallize  at  ordinary  temperatures,  the 
hydrate  FeCl36H20  separates  in  brownish-yellow  crystals. 
This  is  the  ordinary  ferric  chloride  of  commerce.  Several 
other  hydrates  are  known. 

The  oxide  Fe304  occurs  in  nature  in  black  crystals  which 
are  attracted  by  the  magnet.  Some  specimens  are  natural- 
ly magnetic,  and  to  these  the  name  lodestone  is  given. 
Fe304  is  called  magnetic  iron  oxide  or  magnetite,  and  is  an 
important  iron  ore. 

It  is  doubtful  whether  any  salts  corresponding  to  this  oxide 
exist.  The  chloride  would,  of  course,  be  FesCls,  and  it  has  been 
described,  but  it  is  very  likely  only  a  mixture  of  ferric  and  ferrous 
chlorides,  thus  :  FeCl22FeCls. 

377.  Iron  disulphide,  FeS2,  is  the  very  common  mineral 
pyrite.  It  is  brass-yellow,  with  a  metallic  luster,  and  the 
crystals  are  often  cubical.  It  is  too  hard  to  be  scratched 
with  a  knife,  and  hard  enough  to  yield  sparks  when  sharply 
struck  with  a  piece  of  steel.  Great  quantities  of  pyrite  are 
burned  for  the  production  of  sulphuric  acid  (p.  242).  It  is 
of  no  value  as  an  iron  ore,  for  some  of  the  sulphur  would 
find  its  way  into  the  iron  and  reduce  its  tenacity  to  such 
an  extent  as  to  render  it  worthless. 


CHAPTER    XL 

COBALT  AND   NICKEL- THE  PLATINUM  METALS 

378.  Cobalt  and  nickel  resemble  iron,  and  are  remark- 
ably similar  to  each  other.     They  are  both  hard,  white 
metals  with  a  somewhat  gray  luster,  which  melt  only  at  a 
high  white  heat.    They  are  unaffected  by  the  air  at  ordinary 
temperatures,  but  at  a  red  heat  each  becomes  covered  with 
a  scaly  coating  of  the  corresponding  oxide.    Like  iron,  they 
are  attracted  by  the  magnet.     They  are  only  slowly  acted 
upon  by  hydrochloric  and  sulphuric  acids,  but  dissolve  read- 
ily in  nitric  acid.  Chemically,  they  differ  from  iron  chiefly  in 
this  respect,  that  cobalt  and  nickel  each  form  only  one  series 
of  stable  salts — those  in  which  the  metal  is  bivalent.    These 
salts  correspond  to  the  ferrous  compounds.   Simple  salts  cor- 
responding to  the  ferric  compounds  are  unknown  with  both 
nickel  and  cobalt,  though  the  oxides  Co203  and  M203  have 
been  obtained.    Native  nickel  and  cobalt  are  found  only  in 
meteorites,  which  consist,  for  the  most  part,  of  iron  alloyed 
with  rather  large  quantities  of  nickel  and  with  traces  of 
cobalt  and  other  elements.     Cobalt  and  nickel  compounds 
are  generally  found  together  in  nature.    The  chief  ores  are 
the  compounds  of  the  metals  with  sulphur  and  arsenic. 

379.  Cobaltous  oxide,  CoO,  is  a  brown  powder.     This 
oxide   corresponds   to   the   cobalt    salts — CoCl2,Co(NO;!)o, 
CoS04,  and  so  on      These  salts  are  blue  when  free  from 
water,  but  when  they  contain  water  of  crystallization  they 
are  red,   and  their  water   solutions   are   rose-red.      Since 

268 


COBALT  AND  NICKEL  269 

cobalt  nitrate,  cobalt  chloride,  and  the  other  cobalt  salts  all 

have  just  the  same  color  when  dissolved  in  water,  this  must 

++ 

be  the  color  of  the  cobalt  ion,  Co.  On  the  other  hand,  an- 
hydrous cobalt  chloride  is  deep  blue,  and,  when  dissolved 
in  alcohol,  yields  a  blue  solution  which  does  not  conduct  the 
current.  The  blue  color  must  be  due,  therefore,  to  the 
undissociated  molecules  CoCl2.  Owing  to  this  difference 
in  color,  a  water  solution  of  cobalt  chloride  forms  a  kind 
of  sympathetic  ink.  The  color  of  the  cobalt  ion  is  not  very 
intense,  and  writing  made  with  the  dilute  solution  on  paper 
is  invisible  and  remains  invisible  when  the  excess  of  water 
evaporates.  But  when  the  paper  is  warmed,  the  strongly 
colored  anhydrous  salt  is  produced  and  the  writing  appears 
in  blue  characters,  which  again  disappear,  after  a  time,  on 
cooling,  because  the  CoCl2  absorbs  water  from  the  air. 

380.  Uses. — Cobalt  as  metal  has  found  no  commercial 
applications.    Cobaltous  oxide,  when  added  to  melted  glass, 
dissolves  in  it  as  cobalt  silicate,  producing  an  intensely  blue 
mass,  which,  when  cooled  and  powdered,  constitutes  the  blue 
color  called  smalt. 

381.  Nickelotis  oxide,  NiO,  which  is  the  oxide  corre- 
sponding to  the  nickel  salts,  is  green.     Anhydrous  nickel 
salts  are  yellow,  those  containing  water  of  crystallization 

green.     Dilute  water  solutions  of  all  nickel  salts  have  the 

++ 
same  emerald-green  color,  which  is  that  of  the  nickel  ion  Ni. 

382.  Uses  of  Nickel. — On  account  of  its  luster  and  per- 
manence in  the  air,  nickel  is  largely  used  for  plating  arti- 
cles made  of  iron  and  other  metals.     This  is  effected  by 
hanging  the  object  in  a  water  solution  of  a  nickel  salt,  and 
making  it  the  negative  pole  of  a  current  passing  through 
the  liquid.    The  positive  pole  is  a  bar  of  pure  nickel,  which 
dissolves  as  nickel  deposits  at  the  negative  pole,  and  keeps 
the  number  of  nickel  ions  in  the  liquid  the  same.     German 
silver  is  an  alloy  of  nickel  with  copper  and  zinc.    The  nickel 


270  ELEMENTARY  CHEMISTRY 

coins  of  the  United  States,  Germany,  and  Belgium  contain 
25  per  cent  of  nickel  and  75  per  cent  of  copper. 

>.     THE  PLATINUM^  METALS 
joi    I  ,  0^*0 

Ruthenium,  Ru.          Rhodium,  Rh.          Palladium,  Pd. 
Osmium,  Os.  Iridhyn^Ir^  Platinum,  pt^ 

383.  A  reference  to  the  table  of  atomic  weights  will  show  that 
the  first  three  of  these  metals,  Ru,  Rh,  and  Pd,  have  atomic  weights 
in  the  neighborhood  of  100,  while  the  atomic  weights  of  the  last 
three  lie  in  the  neighborhood  of  200.     Here  we  meet  with  a  con- 
firmation of  a  statement  which  has  been  made  before — that  sub- 
stances with  high  atomic  weights  have  high  densities  also,  for  the 
densities  of  the  first  three  are  not  far  from  12,  while  those  of  Os,  Ir, 
and  Pt  are  above  20.      Osmium,  whose  density  is  nearly  22.5,  is  the 
heaviest  of  all  substances.    Palladium  dissolves  in  nitric  acid.    With 
this  exception,   the   platinum  metals  are  unaffected   by  the   three 
ordinary  acids  singly.      Aqua  regia  dissolves  most  of  them,  but 
some,  like  iridium,  are  scarcely  affected  by  it.     They  are  all  white 
metals,  with  a  strong  grayish  metallic  luster,  which  is  permanent  in 
the  air.    They  melt  only  at  high  temperatures,  palladium  having  the 
lowest  melting-point,  while   osmium  has  the  highest.      They  are 
rather  inert  chemically,  and  their  compounds  are  easily  decomposed 
by  heat,  leaving  a  residue  of  the  metal.     They  occur  together  in 
nature  in  rounded  grains  mixed  with  sand,  in  the  same  way  as  the 
placer  deposits  of  gold. 

384.  Platinum  is  by  far  the  most  important  metal  of 
the  group.     It  is  a  white  metal,  with  a  somewhat  grayish 
tint,  capable  of  taking  a  high  polish,  and  very  malleable  and 
ductile.     It  melts  only  at  2000° — a  temperature  much  be- 
yond a  white  heat — but  can  be  melted  in  the  oxyhydrogen 
flame,  and  rapidly  vaporizes  when  subjected  to  the  heat  of 
the  electric  arc.     Finely  divided  platinum  is  black  and  is 
called  platinum  black.    In  this  form  the  metal  best  exhibits 
its  remarkable  power  of  absorbing  gases — a  power  which  is 
shared  by  other  metals  of  the  group,  especially  by  palladium. 
Platinum  black  will  absorb  1,000  times  its  volume  of  oxygen 
or  300  times  its  volume  of  hydrogen.    This  phenomenon  has 


COBALT  AND  NICKEL  271 

engaged  the  attention  of  chemists  for  a  century,  but  we  are 
still  in  doubt  as  to  what  happens.  Either  the  gases  dis- 
solve in  the  metal,  producing  solid  solutions,  or  else  chemi- 
cal combination  occurs  and  unstable  oxides  and  hydrogen 
compounds  of  platinum  are  obtained.  An  important  practi- 
cal point  is,  that  when  platinum  black  is  exposed  to  a  mix- 
ture of  hydrogen  and  oxygen,  both  gases  are  absorbed,  and, 
in  contact  with  the  platinum,  combine  to  form  water  so  rap- 
idly that  the  metal  becomes  red-hot  and  ignites  the  mixture. 
This  is  the  principle  of  the  self-lighting  Welsbach  mantles 
for  gas-burners.  Illuminating  gas  contains  something  like 
half  its  volume  of  hydrogen.  Of  course  the  oxygen  comes 
from  the  air. 

Owing  to  its  high  melting-point  and  its  resistance  to 
the  action  of  most  chemicals,  platinum  is  much  used  for 
the  construction  of  dishes  and  crucibles  which  are  indis- 
pensable in  the  chemical  laboratory.  Platinum  is  not  acted 
upon  by  any  single  acid,  but  is  slowly  dissolved  by  aqua 
regia.  However,  there  are  many  substances  which  do  act 
upon  it,  and  which  must  not  be  heated  in  platinum  vessels. 

385.  Platinum  is  either  bivalent  or  quadrivalent  in  most 
of  its  compounds. 

Platinum  monoxide,  PtO,  is  a  gray,  and  Platinum  dioxide,  PtOi, 
a  black  powder.  Both  are  decomposed  into  platinum  and  oxygen 
by  gentle  heating. 

Platinum  dichloride,  PtCla,  is  a  greenish  powder  insoluble  in 
water. 

Platinum  tetrachloride,  PtCl4,  is  yellow.  It  is  produced 
when  the  metal  is  heated  in  a  current  of  chlorine.  It  is  read- 
ily soluble  in  water,  and,  when  the  liquid  is  evaporated, 
there  separate  red  crystals  of  the  composition  PtCl45H20. 

Both  chlorides  of  platinum  separate  into  metal  and  chlo- 
rine when  strongly  heated. 


CHAPTER    XLI 

CARBON 

0  =  12. 

386.  For  two  reasons  carbon  is  the  most  interesting  of 
all  the  elements.     In  the  first  place,  the  bodies  of  living 
things,  both  animal  and  vegetable,  are  composed  of  carbon 
compounds,  and  the  study  of  these  compounds  has  already 
thrown  much  light,  and  will  throw  much  more,  upon  the 
nature  of  the  chemical  processes  which  are  associated  with 
life.    In  the  second  place,  the  compounds  containing  carbon 
are  marvellously  numerous,  far  more  numerous  than  those 
of  all  the  other  elements  put  together.    We  shall  understand 
clearly  the  reason  for  this  as  we  proceed,  and  shall  perceive, 
at  the  same  time,  that  the  number  of  possible  compounds 
is  infinite. 

Carbon  exists  in  several  different  modifications :  First, 
amorphous  carbon,  of  which  coal  (anthracite),  charcoal, 
and  lampblack  are  more  or  less  impure  varieties.  Second, 
crystalline  carbon,  which  includes  two  very  different  sub- 
stances, graphite  and  the  diamond. 

387.  Coal. — In  the  geological  history  of  the  earth  there 
was  one  epoch,  called  the  Carboniferous  period,  in  which 
vegetation  was  far  more  luxuriant  than  it  is  at  present, 
partly  because  the  temperature  was  higher,  partly  because 
the  air  was  richer  in  carbon  dioxide.     Over  vast  marshy 

272 


CARBON  273 

plains,  not  far  above  sea-level,  grew  forests  of  whose  rich- 
ness and  density  a  tropical  jungle  of  the  present  day  gives 
only  a  faint  idea.  Leaves,  branches,  and  trunks  accumu- 
lated where  they  fell,  and  in  this  way  immense  masses  of 
vegetable  debris  were  produced,  which  were  covered  with 
earth  and  converted  into  coal  by  a  slow  change  in  which  the 
pressure  of  overlying  strata  played  an  important  part. 
Vegetable  matter — wood,  for  instance — contains  mostly  car- 
bon, hydrogen,  and  oxygen,  and  the  change  into  coal  consists 
in  the  removal  of  the  last  two  elements  with  a  portion  of  the 
carbon,  the  rest  of  the  carbon  being  left  with  all  the  mineral 
matter  of  the  original  plant,  as  coal.  This  process  is  very 
gradual,  and  all  stages  of  it  are  found  in  nature.  Here  are 
some  examples : 

1.  Peat  is  plant-substance  only  slightly  carbonized,  and 
still  retaining  a  perfectly  distinct  vegetable  structure. 

2.  Lignite  is  a  brown,  imperfect,  woody  coal  found  in 
the  newer  rocks. 

3.  Bituminous  coal  still  contains  much  hydrogen  and 
some  oxygen.     When  it  is  heated  gaseous  compounds  of 
hydrogen  and  carbon  escape,  which  burn  with  a  bright  flame, 
while  in  the  vessel  there  remains  a  grayish-black  residue, 
called  colce,  composed  of  carbon  with  some  mineral  matter. 
Illuminating  gas  is  made  by  heating  bituminous  coal  in 
fire-clay  retorts  and  collecting  the  gas  given  off,  after  purify- 
ing it,  in  gasometers. 

4.  Anthracite  coal  has  lost  almost  all  of  its  hydrogen, 
and  consists  essentially  of  carbon  with  some  mineral  matter, 
which  remains   as   ash  when  the   coal   is   burned.     When 
heated  in  the  absence  of  air  anthracite  is  almost  unaltered 
and  little  gas  escapes. 

388.  Charcoal  results  when  wood  is  heated  in  the  absence 
of  air.  This  is  best  done  in  iron  retorts,  so  that  the  valuable 
products  which  vaporize  can  be  condensed  and  collected. 
The  properties  of  charcoal  depend  on  the  temperature  at 


274 


ELEMENTARY  CHEMISTRY 


\vhich  it  is  prepared.  When  made  at  a  low  temperature  it 
is  brownish,  and  so  inflammable  that,  if  finely  divided,  it 
catches  fire  on  contact  with  air.  Charcoal  made  at  a  high 
temperature  is  grayish-black,  denser,  and  not  nearly  so  in- 
flammable. 

389.  Lampblack,  or  soot,  is  a  soft  black  powder  which 
is  produced  when  combustible  substances  rich  in  carbon,  like 
turpentine  or  rosin,  are  burned  in  an  insufficient  supply  of 
air.    It  is  employed  in  the  production  of  printer's  ink.    Its 
formation  can  be  illustrated  by  placing  a  cold  porcelain 
plate  in  an  ordinary  luminous  gas  flame,  when  a  coating  of 
lampblack  is  immediately  deposited. 

Crystalline  Carbon 

390.  Graphite  occurs  abundantly,  especially  in  the  older 
rocks.     It  is  soft  and  gray,  with  a  metallic  luster.     When 

drawn  across  paper  it 
leaves  a  black  mark.  A 
mixture  of  graphite  and 
clay  is  employed  for  the 
"  lead "  of  pencils  and 
for  the  manufacture  of 
crucibles,  to  be  employed 
in  the  melting  of  steel 
and  in  other  processes 
requiring  very  high  fem- 

Fie.  51.— Diamond  crystals.  TT77 

peratures.       When    any 

other  modification  of  carbon  is  heated  to  extremely  high 
temperatures  it  is  transformed  into  graphite.  Thus,  when 
a  diamond  is  introduced  into  the  electric  arc  it  swells  up 
and  a  mass  of  graphite  is  produced.  When  a  crucible  made 
of  charcoal  is  heated  in  the  electric  furnace  to  a  tempera- 
ture of  3,000°  or  over  it  is  completely  converted  into  graph- 
ite in  a  few  minutes.  From  facts  of  this  kind  we  know  that, 
at  high  temperatures,  graphite  is  the  stable  modification 


CARBON 


275 


FIG.  52.— Artificial  diamonds. 


of  carbon.  Of  course  one  of  the  modifications  must  be  the 
natural  state  at  ordinary  temperatures,  and  the  others  must 
be  slowly  passing  into  it,  just  as  colorless  phosphorus  slow- 
ly passes  into  the 
red  variety.  But  it 
is  impossible  to 
say  which  modi- 
fication of  carbon 
is  the  stable  form 
at  ordinary  tem- 
peratures, for  the 
speed  with  which 
the  others  change 

into  this  form  (whichever  it  may  be)  is  so  small  that  all 
the  modifications  appear  to  be  permanent. 

391.  Diamond  occurs  native  in  certain  localities,  particu- 
larly in  Brazil  and  at  the  Cape  of  Good  Hope.     It  differs 

from  the  other  forms 
of  the  element  in  its 
high  density  (3.5), 
its  great  hardness, 
and  its  high  index  of 
refraction.  It  exceeds 
all  minerals  in  hard- 
ness, and  one  form  of 
it  called  carbonado, 
which  is  black  and  not 
well  crystallized,  is 
the  hardest  of  all  sub- 
stances. The  diamond 
may  be  colorless,  yel- 
low, brown,  or  black, 
according  to  its  puri- 
ty. Only  the  colorless  varieties  are  esteemed  as  gems — even 
a  very  slight  tint  of  yellow  detracts  immensely  from  the 
19 


PIG.  53.— Preparation  of  diamonds, 
crucible. 


Cooling  the 


276 


ELEMENTARY  CHEMISTRY 


value.    Some  diamond  crystals  are  shown  in  Fig.  51.    The 

diamonds  worn  as  jewels  are  cut  in  such  a  way  as  best  to 

display  the  luster,  and  are  not  crystals. 

Chemically,  the  diamond  is  extremely  inert.     It  is  not 

affected  by  any  acid,  nor,  in  fact,  by  any  liquid  reagent,  and 

most  substances  are 
without  action  upon  it, 
even  at  a  red  heat.  At 
a  bright-red  heat  it 
burns  brilliantly  in 
oxygen  to  carbon  diox- 

FIG.  54a.— Diagram  of  electric  furnace. 

ide.     When  fused  with 

sodium  carbonate  or  potassium  carbonate  it  is  converted 
into  carbon  monoxide,  probably  according  to  the  equation, 


K2C0 


=  2K  +  SCO. 


The  diamond  has  been  obtained  artificially  by  the  French  chemist 
Moissan.  In  order  to  understand  the  method  by  which  this  was  done, 
it  is  necessary,  first  of  all,  to  know  that  carbon  does  not  melt  even  at 
the  very  highest  temperatures.  At  the  heat  of  the  electric  arc  (3500°) 
it  vaporizes  slowly  from  the  solid  state,  and  when  the  vapor  is  con- 
densed by  cooling, 
graphite,  not  dia- 
mond, is  the  prod- 
uct. Now  other  sub- 
stances which  vapor- 
ize from  the  solid 
state  —  arsenic,  for 
instance  —  can  be 
melted  when  heated 
under  pressure.  It 
seemed  to  Moissan 
likely  that,  if  carbon 
could  be  obtained 
liquid  in  a  similar 
way,  diamond  might  be  produced  when  the  liquid  solidified.  But  it 
is  impossible  to  prepare  a  vessel  in  which  carbon  can  be  melted 
under  pressure,  for  no  material  is  known  which  remains  solid  at  the 


FIG.  546. — Electric  furnace  in  operation. 


CARBON  277 

temperature  required.  This  difficulty  was  avoided  in  the  following 
ingenious  way.  When  iron  is  heated  in  contact  with  powdered  char- 
coal in  a  charcoal  crucible  in  the  electric  furnace,  large  quantities  of 
carbon  dissolve  in  the  liquid  metal.  If  the  crucible  is  simply  allowed 
to  cool  in  the  air,  the  carbon  separates  again  as  graphite.  But  if  the 
crucible,  while  at  the  highest  temperature,  is  plunged  into  water, 
the  outside  of  the  mass  of  iron  is  suddenly  cooled  and  a  skin  or  crust 
of  solid  iron  is  formed,  inclosing  the  liquid  interior.  Iron  expands 
when  it  becomes  solid,  and  the  liquid  mass  in  the  inside  in  solidify- 
ing expands  against  the  inclosing  crust  and  produces  an  enormous 
pressure.  Under  this  pressure  the  carbon  separates  from  solution 
in  the  iron,  probably  first  in  minute  drops  of  liquid  carbon,  which 
immediately  solidify  to  diamonds.  The  diamonds  obtained  in  this 
way  are  microscopic  in  size  and  of  no  commercial  value.  Drawings 
of  them  are  shown  in  Fig.  52.  The  cooling  of  the  crucible  is 
shown  in  Fig.  53,  and  the  electric  furnace  in  which  the  heating  is 
conducted  in  Fig.  54.  The  electric  furnace  has  been  employed  by 
Moissan  in  studying  the  effects  of  very  high  temperatures  on  a  great 
variety  of  substances. 


CHAPTEE   XLII 

CARBON  DIOXIDE-CHEMICAL  ENERGY-CARBON  MONOXIDE 
-CARBON  BISULPHIDE 

Two  compounds  of  carbon  with  oxygen  are  known. 

392.  Carbon   dioxide,    C02,   is   produced   when   carbon 
burns  in  an  abundant  supply  of  oxygen  or  air,  and  is  there- 
fore given  off  abundantly  from  fires  in  which  coal  or  wood 
is  burned,  and  from  the  flames  of  gas-jets,  candles,  or  lamps. 
Large  quantities  of  the  gas  are  also  produced  in  the  fermen- 
tation processes  which  are  carried  on  in  breweries,  and  the 
gas  from  this  source  is  liquefied  in  steel  cylinders,  and  in 
this  form  is  largely  used  for  charging  mineral  waters  and 
other  drinks  with  carbon  dioxide,    The  bubbles  which  escape 
from  all  effervescing  drinks  consist  of  carbon  dioxide.    Car- 
bon dioxide  is  produced  by  the  oxidation  of  the  tissues,  and 
leaves  the  body  chiefly  through  the  lungs;  the  expired  air 
contains  about  5  per  cent  of  it  by  volume.    When  either  the 
air  or  the  body  is  in  motion  this  carbon  dioxide  is  rapidly 
blown  away;  otherwise  it  is  only  slowly  removed  by  diffu- 
sion, and  it  has  been  shown  that  the  air  in  the  vicinity  of 
a  man  sitting  quietly  in  a  room  where  there  are  no  air  cur- 
rents may  contain  as  much  as  -J  per  cent  by  volume  of  the  gas. 

393.  Properties  of  carbon  dioxide. — Carbon  dioxide  is  a 
colorless,  odorless  gas  with  a  sharp,  peculiar  taste.    Water 
dissolves  about  one  volume  of  it  at  a  pressure  of  one  atmos- 
phere.   At  higher  pressures  the  amount  dissolved  increases 
almost  in  proportion  to  the  pressure.     We  have  seen  that 
soda-water  is  water  which  has  been  charged  with  the  gas 


CARBON  DIOXIDE  2T9 

under  a  pressure  of  about  four  atmospheres.  Carbon  diox- 
ide is  produced  in  the  alcoholic  fermentation.  From  still 
wines  the  gas  is  all  allowed  to  escape,  the  whole  of  the  fer- 
mentation being  carried  on  in  open  vessels.  In  champagne 
and  other  effervescing  drinks  a  portion  of  the  fermentation 
occurs  in  corked  bottles,  and  the  carbon  dioxide  dissolves  in 
the  liquid  under  its  own  pressure  and  escapes  in  bubbles 
when  the  bottle  is  opened;  hence  the  effervescence. 

Carbon  dioxide  is  easily  converted  into  a  colorless  liquid 
by  pressure  alone ;  when  the  liquid  is  allowed  to  escape 
into  the  air  a  portion  of  it  rapidly  evaporates,  and  this  cools 
the  rest  to  such  an  extent  that  it  freezes  to  a  solid  which 
resembles  snow. 

394.  Action  upon  the  system, — Carbon  dioxide  can  not 
be  classed  as  a  poisonous  gas.     The  workmen  in  the  fer- 
menting cellars  of  breweries  continually  breathe  air  con- 
taining 2  per  cent  or  more  without  damage.    Five  per  cent 
is   injurious,   and   much   more   than  •  that   rapidly   causes 
death  by  suffocation.    When  the  body  is  plunged  in  a  vessel 
containing  carbon  dioxide,  the  head  being  left  free  so  that 
pure  air  is  respired,  there  is  at  first  a  tingling  sensation  of 
warmth  over  the  skin.     This  is  followed  by  such  alarming 
symptoms  of  collapse  that  the  experiment  must  be  discon- 
tinued. 

395.  Importance  of  carbon  dioxide  in  the  life  process. — 
When  we  studied  the  atmosphere  we  learned  that  the  air 
contains  about  3  parts  in  10,000  of  carbon  dioxide,  and  that 
without  this,  trifling  as  it  seems,  life  upon  our  planet  would 
be  impossible.    In  sunlight,  in  the  green  parts  of  plants,  the 
carbon  dioxide  is  decomposed,  the  oxygen  returned  to  the 
air,  and  the  carbon,  with  hydrogen  and  oxygen  from  water 
and  nitrogen  from  other  sources   (p.  101),  built  up  into 
complicated  compounds  of  these  four  elements  which  serve, 
directly  or  indirectly,  as  food  for  all  animals  and  plants. 
This  may  well  be  called  the  most  important  of  all  chemical 


280  ELEMENTARY  CHEMISTRY 

changes.  Its  importance  is  due  to  the  fact  that  it  is  the 
only  place  where  energy  enters  the  life  process.  In  this 
change  the  energy  of  sunlight  is  absorbed;  the  compounds 
which  the  plant  makes  contain  more  energy  than  the  mate- 
rials from  which  it  makes  them.  Every  other  process  which 
occurs  in  a  plant,  and  every  change  without  exception  in  the 
animal,  is  associated  with  a  loss  or  dissipation  of  energy 
which  appears  (in  the  animal,  especially)  as  motion  and 
heat. 

396.  Carbonic  acid. — The  solution  of  carbon  dioxide  in 
water  is  acid.     This  is  due  to  hydrogen  ions  derived  from 
the  dissociation  of  the  acid,  H2C03,  which  is  produced  thus : 

H20  +  C02  =  H2C03. 

H2C03  is  called  carbonic  acid.  It  is  highly  unstable, 
and  is  only  known  in  solution,  but  its  salts,  called  carbonates, 
are  stable  and  important.  Some  of  them  we  have  already 
studied.  They  all,  when  treated  with  hydrochloric  or  nitric 
acid,  give  off  carbon  dioxide  with  effervescence,  and  in  this 
way  we  can  tell  whether  a  given  substance  is  a  carbonate 
or  not.  The  action  of  hydrochloric  acid  upon  calcium  car- 
bonate is  the  best  method  of  making  carbon  dioxide.  Small 
lumps  of  marble  are  placed  in  an  ordinary  gas-generating 
bottle,  water  enough  to  cover  them  is  poured  through  the 
funnel-tube,  and  then  strong  hydrochloric  acid — 
CaC03  +  2HC1  =  CaCl2  +  H20  +  C02. 

The  gas  is  collected  by  downward  displacement. 

397.  Chemical  energy. — It  is  a  familiar  fact  that  when 
charcoal  or  coal  is  burned  to  carbon  dioxide,  large  quantities 
of  heat  and  light  are  produced.     Now,  heat  and  light  are 
forms  of  energy,  and  it  is  clear  that  charcoal  and  oxygen 
separately  must  contain  more  energy  than  after  they  have 
combined  to  carbon  dioxide.     Sodium  burns  brilliantly  in 
chlorine  to  salt  (p.  64),  and  it  follows  from  this  that  salt, 
like  carbon  dioxide,  must  contain  less  energy  than  its  con- 


CHEMICAL  ENERGY  281 

stituents  do  when  they  exist  as  separate  elements.  If  we 
desire  to  separate  the  salt  into  sodium  and  chlorine  again, 
all  of  the  energy  which  was  given  off  during  the  combina- 
tion must  be  restored,  for  energy,  like  matter,  can  not  be 
created,  and  if  we  could  decompose  the  salt  without  sup- 
plying energy  to  it,  the  energy  which  the  sodium  and 
chlorine  would  then  contain  would  have  been  obtained 
without  any  expenditure — made  out  of  nothing.  By  re- 
peating the  process,  we  could  manufacture  any  desired 
quantity  of  energy,  which  is  impossible.1  The  most  con- 
venient way  of  supplying  energy  to  the  salt  is  to  melt  it 
and  pass  through  it  an  electric  current,  when  it  absorbs 
the  energy  of  the  current  and  separates  into  sodium  and 
chlorine. 

Most  compounds  are  like  salt  and  carbon  dioxide.  They 
contain  less  energy  than  their  constituents.  Hence  chemi- 
cal combination  is  usually  attended  by  the  evolution  of 
heat,  or  of  light,  or  of  both.  However,  there  are  many  ex- 
ceptions. Many  compounds — nitrogen  chloride  (p.  117)  is 
a  good  example — contain  more  energy  than  their  constitu- 
ents taken  separately.  Such  compounds  are  frequently  ex- 
plosive, because  they  tend  to  separate  into  their  elements, 
at  the  same  time  liberating  energy.  The  familiar  gas  acety- 
lene is  another  example.  It  is  a  compound  of  carbon  and 
hydrogen,  C2H2,  and  when  the  two  elements  unite  to  form 
it,  energy  is  absorbed.  Hence  it  tends  to  separate  again 
with  explosion,  and  until  this  fact  was  learned  and  proper 
care  was  applied  in  dealing  with  it,  accidents  with  it  were 
frequent. 

In  ordinary  life  it  is  often  necessary  to  store  up 
energy.  Thus,  in  winding  a  clock  run  by  weights,  we  raise 

1  Show  by  the  same  argument  that  when  carbon  is  first  burned  to 
carbon  monoxide  and  then  the  carbon  monoxide  to  carbon  dioxide, 
the  same  quantity  of  energy  must  be  liberated  as  though  the  carbon 
had  been  burned  at  once  to  carbon  dioxide. 


282  ELEMENTARY  CHEMISTRY 

the  weights  and  store  up  energy  which  suffices  to  keep  the 
mechanism  in  motion  for  days.  In  the  same  way  electrical 
energy  can  be  stored  up  for  a  time — in  Leyden  jars,  for  in- 
stance. Evidently  chemical  energy  can  be  kept  in  the  same 
way.  We  might,  for  instance,  apply  energy  to  the  decom- 
position of  salt,  and  then  preserve  the  sodium  and  chlorine 
separately.  At  any  desired  future  time  the  energy  could 
be  obtained  again  by  simply  bringing  the  two  together. 
But  it  would  be  better  to  use  carbon  dioxide  instead  of  salt, 
for  it  would  not  be  necessary  to  preserve  the  oxygen,  since 
the  air  furnishes  i  plentiful  supply  of  this  constituent.  It 
would  be  quite  sufficient  to  lay  aside  the  carbon,  and,  at 
any  time,  the  energy  spent  in  decomposing  the  carbon  diox- 
ide could  all  be  recovered  by  simply  burning  the  carbon  in 
the  air. 

Chemical  energy  is  the  great  form  in  which  energy  is 
stored  away  for  practical  purposes.  It  has  this  great  advan- 
tage, that  it  can  be  preserved  for  any  length  of  time  with- 
out much  loss,  while  light,  heat,  and  other  forms  of  energy 
are  rapidly  dissipated.  The  chief  substances  employed  as 
storehouses  of  chemical  energy  are,  first,  impure  forms  of 
carbon  (charcoal,  coke,  and  coal),  and  second,  various  com- 
pounds and  mixtures  containing  carbon  and  hydrogen — 
wood  and  petroleum,  for  instance.  It  is  not  necessary  for 
us  to  start  with  carbon  dioxide  and  supply  energy  to  it  tf> 
decompose  it,  for  this  has  been  done  for  us  by  the  plants 
of  the  coal-forming  period  (p.  272),  which  absorbed  the 
energy  of  sunlight  and  stored  it  up  as  chemical  energy. 
This  is  a  remarkable  instance  of  the  permanence  of  chemi- 
cal energy.  Whether  we  burn  coal  or  wood,  the  energy  which 
appears  is  derived  from  sunlight.  The  only  difference  is 
that  in  the  case  of  wood  the  energy  was  stored  up  quite 
recently,  while  the  energy  of  coal  is  derived  from  the  sun- 
shine which  fell  upon  plants  which  grew  hundreds  of  thou- 
sands of  years  ago. 


CARBON  MONOXIDE  283 

398.  Carbon  monoxide,  CO,  is  produced  when  carbon 
dioxide  comes  into  contact  with  red-hot  carbon.  When  a 
slow  stream  of  carbon  dioxide  is  passed  through  a  wide 
glass  tube  about  60  cm.  long  full  of  lumps  of  charcoal  and 


FIG.  55.— Production  of  carbon  monoxide  from  carbon  dioxide  and  carbon. 

heated  to  redness  in  a  furnace,  almost  pure  carbon  monoxide 
issues  from  the  tube  (Fig.  55). 

C02  +  C  =  2CO. 

This  reaction  occurs  in  an  ordinary  coal  fire.  The  oxy- 
gen which  enters  below,  at  the  grate,  combines  with  the 
carbon  to  C02.  This,  as  it  rises  through  the  glowing  coal 
above,  is  converted  into  carbon  monoxide,  which  finally, 
when  it  reaches  the  top,  burns  again  to  carbon  dioxide  at 
the  expense  of  the  oxygen  of  the  air,  producing  the  blue 
flame  which  plays  upon  the  surface  of  the  coal. 

Carbon  monoxide  is  usually  made  in  the  laboratory  by  heating 
oxalic  acid,  H3CaO4,  with  strong  sulphuric  acid. 
H2C3O4  =  H,0  -|-  CO,  +  CO. 

The  water  produced  is  retained  by  the  sulphuric  acid.  A  mix- 
ture of  equal  volumes  of  CO  and  COa  escapes,  and  this  is  passed 
through  a  wash-bottle  containing  potassium  hydroxide,  KOH,  which 
absorbs  the  COa.  The  CO  passes  on  and  is  collected  over  water 
(Fig.  56). 


284  ELEMENTARY  CHEMISTRY 

399.  Properties. — Carbon  monoxide  is  a  colorless,  odor- 
less gas,  very  slightly  soluble  in  water  and  difficult  to 
liquefy.  In  air  or  oxygen  it  burns  with  a  pure  blue  flame 


FIG.  56. — Preparation  of  carbon  monoxide. 

to  carbon  dioxide.  It  is  very  poisonous,  |  per  cent  in 
air  being  rapidly  fatal  to  all  animals,  and  much  less  than 
that  if  the  inhalation  is  prolonged.  This  is  due  to  the 
fact  that  it  combines  with  the  haemoglobin  of  the  red 
blood-corpuscles,  producing  a  stable  compound  and  pre- 
venting them  from  doing  their  work  of  carrying  oxygen 
about  the  body.  Death  by  carbon  monoxide  poisoning  is 
therefore  nothing  but  suffocation  brought  about  in  an  un- 
usual way. 

If  an  animal  is  put  in  air  under  a  pressure  of  ten  atmospheres,  as 
much  as  6  per  cent  of  carbon  monoxide  can  be  mixed  with  the  air 
without  causing  any  symptoms  of  poisoning.  This  is  because,  at 
the  high  pressure,  enough  oxygen  dissolves  in  the  plasma  of  the 
blood  to  supply  the  needs  of  the  tissues,  so  that  the  fact  that  the 
corpuscles  no  longer  act  as  oxygen-carriers  makes  no  difference. 

400.  Considerable  quantities  of  carbon  monoxide  are 
contained  in  coal  gas,  and  still  larger  quantities  (40  per 


CARBON   DISULPHIDE 


285 


cent  or  more)  in  water  gas,  which  is  made  by  injecting 
steam  into  a  mass  of  red-hot  coal, 

H20  +  C  =  CO  +  H2, 

and  which  is  largely  mixed  with  coal  gas  at  present  in  the 
illuminating  gas  supplied  to  most  cities.  Water  gas,  as 
the  equation  shows,  is  a  mixture  of  carbon  monoxide  and 
hydrogen. 

401.  Carbon  disulphide,  CS2,  is  made  by  throwing  sul- 
phur into  a  cast-iron  cylinder  full  of  charcoal  heated  to  red- 
ness (Fig.  57).  It  is  liberated  as  vapor  which  is  condensed 
by  cooling.  It  is  a  colorless  liquid,,  denser  than  water,  in 
which  it  is  almost  insoluble.  When  pure  its  odor  is  faint 
and  rather  pleasant,  but  it  usually  contains  impurities  which 
impart  to  it  a  disgusting  smell.  It  inflames  with  extreme 


FIG.  57.— Manufacture  of  carbon  disulphide.    A,  cylinder  of  charcoal ;  J,  emp 
vessel  in  which  the  CSa  partly  condenses  ;   T,  condenser. 


readiness  and  burns  to  carbon  dioxide  and  sulphur  dioxide. 
Its  vapor  is  somewhat  poisonous,  and  the  workmen  engaged 
in  its  manufacture  suffer  from  disorders  of  the  nervous 
system. 

f 


CHAPTER   XLIII 


SOME   CARBON  COMPOUNDS 

402.  Methane,  CH4. — When  the  mud  at  the  bottom  of  a 
swampy  pool  is  stirred,  bubbles  of  gas  rise  to  the  surface., 
With  the  help  of  an  inverted  bottle  filled  with  water  and  a 
funnel  to  catch  the  bubbles,,  it  is  easy  to  collect  some  of  the 

gas  (Fig.  58).  In  this 
way  we  can  ascertain  that 
it  is  colorless  and  combus- 
tible, burning  with  a  pale 
yellow  flame.  It  is  meth- 
ane, CH4,  the  simplest 
compound  of  carbon  and 
hydrogen.  It  is  also  called 
marsh  gas. 

Methane  can  be  made 
artificially  in  various 
ways,  one  of  which  is  to 
mix  hydrogen  sulphide, 
H2S,  with  the  vapor  of 
carbon  disulphide,  CS2,  and  pass  the  mixture  over  red-hot 
copper.  The  copper  combines  with  the  sulphur,  forming 
Cu2S,  and  the  carbon  and  hydrogen  unite,  thus : 

8Cu  +  CS2  +  2H2S  =  4Cu2S  +  CH4. 

When  methane  is  burned,  the  carbon  is  converted  into 
carbon  dioxide  and  the  hydrogen  to  water.    The  same  two 
products  are  obtained  when  other  compounds  containing 
286 


Fie.  58a.— Collecting  marsh  gas. 


SOME  CARBOtf  COMPOUNDS  287 

carbon  and  hydrogen  are  burned.    The  chief  results  of  the 
burning  of  a  candle,  an  oil  lamp,  or  a  wood  fire  are,  there- 
fore, carbon  dioxide  and  water.    Anthracite  coal,  which  con- 
tains  almost   no   hydrogen, 
produces    chiefly    carbon 
dioxide.     The  ash  which  is 
obtained  by  the  burning  of 
coal  is  simply  incombustible 
mineral  matter  which  exist- 
ed   originally    in    the    fuel, 
and  is  not  a  product  of  the 
combustion. 

The  slow  conversion  of 

Vegetable    matter    into    COal         Fio.  58*. -Transferring  marsh  gas  t 

has    been    referred    to     (p. 

273).  Methane  is  produced  in  this  process,  and  it  fre- 
quently escapes  in  large  quantities  from  beds  of  bituminous 
coal.  Mingling  with  the  air  of  the  mine  it  forms  the  ex- 
tremely dangerous  explosive  mixture  which  miners  call 
fire-damp?  The  carbon  dioxide  produced  by  the  explosion 
is  suffocating,  and  forms  the  "  choke-damp "  or  "  after- 
damp "  which  is  likely  to  suffocate  those  whom  the  ex- 
plosion has  left  alive,  and  is  especially  dangerous  to  res- 
cuing parties. 

1  Like  methane,  every  combustible  gas  and  vapor  forms  an  explo- 
sive mixture  with  air.  Thus,  if  si  cylinder  is  filled  with  illuminating 
gas  and  lighted,  it  burns  quietly,  because  it  can  only  burn  at  the  sur- 
face of  separation  between  the  gas  and  air;  but  if  the  gas  is  mixed 
with  the  proper  quantity  of  air,  it  explodes  when  a  flame  is  applied 
because  the  combustion  extends  instantly  through  the  whole  mass. 
For  this  reason  substances  which  give  off  combustible  vapors,  like 
ether,  gasoline,  and  naphtha,  should  never  be  used  when  a  flame 
or  a  fire  is  anywhere  in  the  neighborhood.  Even  combustible 
solids,  like  coal-dust  and  flour,  when  suspended  in  air  produce  ex- 
plosive mixtures.  Destructive  explosions  have  occurred  in  flour- 
mills  from  sparks  from  the  milling  stones  igniting  the  mixture  of 
flour  and  air. 


288  ELEMENTARY  CHEMISTRY 

403.  Ethane,  C2H6,  is  a  colorless,  combustible  gas  more 
easily  condensed  to  a  liquid  than  methane. 

Methane  and  ethane  are  the  first  two  members  of  a  long 
series  of  compounds  of  carbon  and  hydrogen.  The  series  has 
this  peculiarity,  that  in  every  member  of  it  the  number  of  hy- 
drogen atoms  in  the  formula  can  be  obtained  by  multiplying 
the  number  of  carbon  atoms  by  two  and  adding  two  to  the 
product.  Thus  in  ethane  the  number  of  carbon  atoms  is  2  and 
the  number  of  hydrogen  atoms  is  2  X  2  +  2  =  6.  Compounds 
of  carbon  and  hydrogen  are  called  hydrocarbons.  What  is  the 
formula  of  the  hydrocarbon  of  the  methane  series  which  con- 
tains 9  carbon  atoms?  The  number  of  hydrogen  atoms 
must  be  9X2  +  2  or  20 ;  hence  the  formula  must  be 
C9H20.  This  hydrocarbon  is  a  liquid  called  nonane.  Ordi- 
nary kerosene  or  burning  oil  consists  largely  of  it.  The 
higher  members  of  the  series  are  white,  odorless  solids  re- 
sembling paraffine,  which  is  a  mixture  containing  a  number 
of  them.  The  highest  member  which  has  been  made  so  far 
is  C60H122.  Crude  petroleum  is  a  mixture  containing  many 
of  the  hydrocarbons  of  this  series.  By  distillation  it  is  sepa- 
rated into  various  fractions,  like  naphtha,  benzine,  burning 
oil,  and  lubricating  oil,  all  of  which  have  important  uses. 

404.  Other  series  of  hydrocarbons  are  Tcnown.     Acety* 
lene,  C2H2,  is  the  first  member  of  a  series  in  which  the 
number  of  hydrogen  atoms  is  obtained  by  multiplying  the 
number  of  carbon  atoms  by  two  and  subtracting  two  from 
the  product.     It  is  interesting  because  it  can  be  made  in 
small  quantities  by  allowing  an  electric  arc  to  burn  between 
carbon  poles  in  an  atmosphere  of  hydrogen.     For  years  it 
was  supposed  to  be  the  only  hydrocarbon  which  could  be 
obtained  by  the  direct  union  of  its  elements,  but  very  re- 
cently it  has  been  proved  that  traces  of  methane  and  ethane 
can  be  made  in  the  same  way.    Acetylene  is  a  colorless  gas 
with  a  peculiar  odor.     Its  mixture  with  air  is  violently 
explosive,  more  so  than  is  the  case  with  other  combustible 


SOME  CARBON  COMPOUNDS 

gases,  so  that  more  care  is  required  in  working  with  it.  It 
is  considerably  used  as  an  illuminant,  especially  in  small 
places,  where  it  would  not  pay  to  erect  a  large  gas  works, 
for  its  manufacture  requires  little  apparatus.  Acetylene  is 
prepared  on  the  large  scale  and  in  the  laboratory  by  the  ac- 
tion of  water  upon  calcium  carbide,  CaC2,  calcium  hydroxide 
(slaked  lime)  being  produced  at  the  same  time. 
CaC2  +  2H20  =  Ca(OH)2  +  C2H2. 

405.  Calcium   carbide,    CaC2,   forms   colorless    crystals 
when  pure,  but  the  commercial  product  is  a  hard  iron- 
black  mass.    When  it  comes  into  contact  with  water  acetylene 
is  liberated. 

The  reaction  is  violent,  and,  on  the  large  scale,  the  appa- 
ratus for  preparing  acetylene  is  so  constructed  that  pow- 
dered calcium  carbide  is  gradually  fed  into  water.  In  this 
way  a  regular  stream  of  gas  is  obtained  and  the  reaction 
can  be  readily  controlled.  Calcium  carbide  is  made  by  heat- 
ing a  mixture  of  lime  and  charcoal  to  about  3,000°  by  means 
of  an  electric  arc  passing  through  the  mass. 

406.  Substitution  products  of  the  hydrocarbons. — When 
chlorine  acts  upon  methane  a  hydrogen  atom  is  driven  out 
and  a  chlorine  atom  takes  its  place  in  the  molecule,  forming 
a  compound  whose  formula  is  CH3C1.    The  hydrogen,  how- 
ever, is  not  liberated.    It  combines  with  more  chlorine,  pro- 
ducing hydrochloric  acid,  so  that  the  change  is  described 
by  the  equation — 

CH4  +  C12  =  CH3C1  +  HC1. 

By  the  further  action  of  chlorine  the  other  three  hydro- 
gen atoms  can  one  by  one  be  substituted  by  chlorine,  so  that 
by  the  action  of  chlorine  upon  methane  the  following  four 
compounds  can  be  obtained : 
CH3C1, 
CH2C12, 

CHC13,  and  finally 
CC14. 


290  ELEMENTARY  CHEMISTRY 

Since  the  compound  CC14  contains  no  hydrogen  to  be 
replaced,  chlorine  has  no  action  upon  it.  Similar  com- 
pounds containing  bromine  and  iodine  can  be  made,  and  the 
action  of  chlorine  upon  other  hydrocarbons  is  similar. 
Thus  by  the  prolonged  action  of  chlorine  upon  ethane,  C2H6, 
all  the  hydrogen  can  be  gradually  replaced  by  chlorine,  pro- 
ducing, first,  C2H5C1,  and  finally  C2C16. 

Such  compounds  are  called  substitution  products.  Only 
two  of  them  are  sufficiently  important  to  be  mentioned  here. 

407.  Chloroform,  CHC13,  results  when  three  of  the  hy- 
drogen atoms  of  methane  are  substituted  by  chlorine.    This 
is  not  the  practical  method  of  making  it.    It  is  prepared  by 
a  somewhat  complicated  reaction  which  takes  place  when 
alcohol  is  distilled  with  bleaching  powder  (p.  88).     It  is  a 
colorless  liquid  with  a  peculiar  fragrant  odor.    It  is  largely 
used  for  producing  insensibility  in  surgical  operations, 

408.  lodoform,  CHI3,  is  the  corresponding  iodine  com- 
pound.   It  is  a  yellow  crystalline  powder  with  an  offensive 
smell.    It  is  much  employed  in  surgery  for  antiseptic  dress- 
ings for  wounds.    Such  a  dressing  prevents  the  entrance  of 
germs  from  the  outside  into  a  wound,  and  allows  it  to  heal 
without  the  formation  of  pus  or  other  disturbing  compli- 
cations.    The  importance  of  this  treatment  is   immense. 
Every  year  thousands  of  lives  are  saved  by  it. 

409.  Alcohols. — We  have  just  seen  that  the  hydrogen  of 
hydrocarbons  can  be  replaced  by  chlorine  and  other  elements. 
It  can  also  be  replaced  by  groups  or  radicals,  like  OH  (pp. 
109-110).    The  compounds  produced  by  the  substitution  of 
hydroxyl  for  one  or  more  of  the  hydrogen  atoms  of  a  hydro- 
carbon are  called  alcohols.    For  instance,  by  processes  which 
we  need  not  describe  one  hydrogen  atom  of  methane  can  be 
replaced  by  hydroxyl,  OH.     The  resulting  substance  must 
have  the  formula  CHSOH.    It  is  the  simplest  alcohol. 

410.  Methyl  alcohol,  CH3OH,  is  also  called  wood  alco- 
hol, because  it  is  produced  when  wood  is  distilled  for  the 


SOME  CARBON  COMPOUNDS  291 

production  of  charcoal,  and  this  is  the  industrial  method  of 
obtaining  it.  It  is  a  colorless  liquid  with  an  alcoholic  odor. 
It  dissolves  many  resinous  substances,  shellac,  for  instance, 
and  is  therefore  employed  in  the  manufacture  of  certain 
kinds  of  varnish.  When  a  wooden  surface  is  coated  with 
a  solution  of  shellac  in  wood  alcohol,  the  alcohol  rapidly 
evaporates  and  leaves  the  shellac  as  a  lustrous  film. 

Methyl  alcohol  readily  takes  fire  and  burns  with  a  hot 
blue  flame  to  carbon  dioxide  and  water.  When  its  vapor 
is  mixed  with  air  and  led  over  a  hot  spiral  of  platinum  wire 
a  kind  of  incomplete  combustion  takes  place.  Two  hydro- 
gen atoms  are  removed  from  the  CH40  and  converted  into 
water,  while  a  compound,  CH20,  called  formaldehyde,  is  left. 

CH40  +  0  =  CH20  +  H20. 

411.  Formaldehyde,  CH20,  is  a  gas  which  can  be  con- 
densed to  a  colorless  liquid.    Its  odor  is  pungent  and  suffo- 
cating.   It  is  rapidly  fatal  to  bacteria,  and  is  the  most  con- 
venient and  the  most  frequently  employed  of  all  disinfec- 
tants.    For  this  purpose  lamps  are  constructed  in  which 
formaldehyde  is  vaporized  by  heat,  and  the  vapor  escapes 
through  a  narrow  tube  which  can  be  passed  through  the 
keyhole  of  the  room  to  be  disinfected. 

412.  Ethyl  alcohol,  C2H5OH,  is  the  substance  ordinarily 
called  alcohol.     It  is  a  colorless  liquid  which  freezes  at 
—130°  to  a  white  mass.     Alcohol  burns  easily.     The  flame 
is  blue  and  almost  non-luminous,  and  the  products  are  car- 
bon dioxide  and  water. 

413.  Fermentation. — The  alcohol  and  alcoholic  drinks  of 
commerce  are  produced  by  the  fermentation  of  liquids  con- 
taining glucose,  C6H1206,  a  substance  strongly  resembling 
ordinary  sugar  in  composition  and  properties.    In  the  pres- 
ence of  the  yeast-plant  it  separates  into  alcohol  and  carbon 
dioxide — 

C6H1206  =  2C2H5OH  +  2C02. 


292  ELEMENTARY  CHEMISTRY 

Thus,  the  fresh  juice  of  apples  or  grapes  is  sweet  on 
account  of  the  glucose  which  it  contains.  On  being  allowed 
to  stand,  it  ferments,  carbon  dioxide  escapes,  and  the  sweet 
taste  is  replaced  by  the  alcoholic  flavor  of  wine  or  cider,  as 
the  case  may  be.  Fermentation  is  not  produced  by  any 
vital  process  of  the  yeast  plant.  Yeast  is  killed  by  being 
soaked  for  a  short  time  in  a  mixture  of  alcohol  and  ether, 
and  the  dead  yeast  so  obtained  on  being  thrown  into  a 
glucose  solution  produces  a  violent  fermentation.  When 
a  mass  of  yeast  cells  is  ground  for  a  time  with  fine  angular 
sand  and  water,  so  as  to  crush  the  cells  and  let  their  con- 
tents escape,  and  the  pasty  mixture  is  subjected  to  strong 

pressure,  a  liquid  escapes  which 
can  be  completely  freed  from 
yeast  cells  by  filtering.  This 
liquid,  when  added  to  a  solution 
of  glucose  or  ordinary  sugar, 
throws  it  into  energetic  fer- 
mentation. 

Facts  of  this  kind  show  that 
what  the  yeast  plant  does  is  to 
produce  and  store  up,  in  the  in- 
terior of  the  cells,  a  substance 
which  is  able  to  act  catalytically 
(pp.  42  and  49)  upon  the  glu- 
cose and  separate  it  into  alcohol 

Pio.  59a.-Single  yeast  cell.  *^   Cai>b°n   di°xide'       The   nam6 

zymase  has   been   given   to   this 

substance.  It  is  white  and  soluble  in  water,  but  has  not  yet 
been  obtained  in  anything  like  pure  condition.  The  yeast 
plant  is  shown  in  Fig.  59,  greatly  enlarged. 

Details  regarding  the  production  of  alcoholic  beverages 
must  be  looked  for  in  special  works.  It  may  be  remarked 
that  malt  liquors — beer,  ale,  and  stout — contain  from  2  to 
10  per  cent  of  alcohol,  and  wines  a  somewhat  larger  quan- 


SOME  CAUBON  COMPOUNDS  293 

tity,  up  to  20  per  cent  in  port.    Spirituous  liquors — whisky, 
brandy,  etc. — contain  50  per  cent  or  more. 

414.  In  order  to  obtain  purer  alcohol  from  fermented 
liquids,  distillation  is  resorted  to.  The  alcohol  boils  at  a 
lower  temperature  than  the  water  and  other  substances  pres- 


FIG.  595.— Yeast  in  various  stages  of  development. 

ent,  and  so  its  vapor  passes  off  more  readily.  Distilling 
apparatus  of  great  perfection  has  been  devised  for  this 
purpose,  and,  by  means  of  it,  it  is  possible  to  obtain  a  prod- 
uct containing  less  than  5  per  cent  of  water.  This  is  suffi- 
cient for  commercial  uses,  but  frequently,  in  the  laboratory, 
alcohol  quite  free  from  water  is  required.  This  can  be 
obtained  by  letting  95  per  cent  alcohol  stand  in  contact 
with  lumps  of  lime,  which  gradually  slakes  and  absorbs  the 
water.  Then  when  the  alcohol  is  distilled  it  is  almost 
water-free.  The  last  traces  of  water  can  be  removed  by 
treating  the  alcohol  with  small  pieces  of  sodium  and  dis- 
tilling again.  Anhydrous  alcohol  is  commonly  called  abso- 
lute alcohol.  It  absorbs  moisture  from  the  air  and  must 
be  kept  in  well-closed  vessels.  It  is  quite  poisonous,  partly 
owing  to  the  energy  with  which  it  takes  water  from  the 
tissues  of  the  body. 


294:  ELEMENTARY  CHEMISTRY 

415.  Aldehyde,  C2H40,  is  made  by  the  gentle  oxidation 
of  alcohol — 

C2H60  +  0  =  C2H40  +  H20, 

just  as  formaldehyde  is  made  from  methyl  alcohol.  It  is 
a  colorless  liquid  with  an  odor  suggesting  that  of  apples. 
It  has  no  application,  but  a  compound  in  which  three  of  its 
hydrogen  atoms  are  replaced  by  chlorine  is  important  in 
medicine.  It  is  called 

416.  Chloral,  C2HC130. — It  crystallizes  with  one  mole- 
cule of  water.     In  small  doses  it  produces  sleep,  in  large 
doses  insensibility  or  even  death. 

417.  Acetic  acid,  C2H402. — We  have  seen  that  aldehyde 
C2H40  is  produced  by  the  oxidation  of  alcohol.    Now  when 
aldehyde  is  dissolved  in  water  and  the  solution  exposed  to 
the  air,  the  peculiar  odor  of  the  aldehyde  disappears,  and, 
after  a  time,  the  liquid  smells  and  tastes  like  vinegar.     At 
the  same  time  it  becomes  acid  to  litmus  and  other  indicators. 
In  this  process  an  atom  of  oxygen  is  taken  from  the  air  by 
each  molecule  of  aldehyde  and  acetic  acid,  C2H402,  is  pro- 
duced.   Acetic  acid  can  be  obtained — without  making  alde- 
hyde— by  the  vigorous  oxidation  of  alcohol — 

C2H60  +  02  =  C2H402  +  H20. 

Acetic  acid  is,  at  low  temperatures,  a  colorless  crystalline 
solid  which  melts  a  little  below  room-temperature  to  a  color- 
less liquid  with  a  pungent  smell.  It  mixes  with  water  in  all 
proportions.  It  is  produced  when  wood  is  distilled,  and 
this  is  the  commercial  source  of  the  pure  acid. 

418.  Vinegar  is  a  dilute  water  solution  containing  5  to 
15  per  cent  of  acetic  acid.    It  also  contains  various  flavoring 
and  coloring  constituents  which  differ  in  different  vinegars, 
according  to  the  mode  of  preparation.    Any  dilute  alcoholic 
liquid — any  wine  or  cider,  for  instance — will  change  to  vine- 
gar if  exposed  to  the  air.     The  change  of  the  alcohol  into 


SOME  CARBON  COMPOUNDS 


295 


FIG.  60.  — Quick  vinegar  process. 
K,  cask ;  T7,  thermometer ;  L, 
perforated  bottom  ;  Z  Z  Z,  aper- 
tures for  admission  of  air;  Et 
reservoir  for  collection  of  vine- 
gar. 


acetic  acid  is  brought  about  by  a  peculiar  micro-organism 

called  bacterium  aceti,  which  exists  in  the  liquid.     The 

equation    above    shows    that    a 

molecule  of  oxygen  from  the  air 

is  required  for  each  molecule  of 

alcohol.     This  explains  the  fact 

that  the  bung  is  removed  from 

a   barrel   of   cider   when    it   is 

desired  to  allow  the  liquid  to 

turn  to  vinegar.     In  the  quick 

vinegar    process    shavings    are 

placed    in    a    large    perforated 

cask  and  are  wet  with  vinegar 

to   supply   the    micro-organism. 

Wine,   cider,   or   dilute    alcohol 

is   allowed  to   trickle   over  the 

shavings   (Fig.  60).     Thus  the 

surface   exposed   to   the   air   is 

made  very  great,  and  the  production  of  vinegar  becomes 

rapid. 

419.  Like  ordinary  alcohol,  methyl  alcohol  yields  an 
acid  when  oxidized — 

CH40  +  02  =  CH202  +  H20. 

Formic  acid,  CH202,  is,  like  acetic  acid,  a  colorless 
liquid  with  a  pungent  odor.  Its  name  is  derived  from  the 
fact  that  it  exists  in  ants. 

Formic  and  acetic  acids  are  weak  acids — that  is,  they  are 
much  less  separated  into  ions  when  dissolved  in  water  than 
are  hydrochloric  or  nitric  acid.  Hence  the  acid  properties, 
which  are  due  to  the  presence  of  hydrogen  ions,  are  feeble. 

420.  Series  of  carbon  compounds. — Wood  alcohol   and 
ordinary  alcohol  are  the  first  two  members  of  a  series  of 
alcohols  which  corresponds  to  the  series  of  hydrocarbons,  of 
which  methane  stands  at  the  head.    An  alcohol  correspond- 
ing to  every  hydrocarbon  can  be  prepared.    Those  alcohols 


296  ELEMENTARY  CHEMISTRY 

standing  just  above  ordinary  alcohol  in  the  series  are  liquids 
with  the  odor  of  "  fusel  oil"  which  is  a  mixture  of  several 
of  them.  The  higher  members,  like  C30H620,  are  solids 
resembling  paraffine. 

There  is  also  a  series  of  aldehydes  in  which  one  member 
corresponds  to  each  of  the  methane  hydrocarbons  and  a 
similar  series  of  acids.  In  both  those  series  also  the  higher 
members  are  waxy  solids. 

If  the  student  will  reflect  that  hundreds  of  different 
hydrocarbons  are  known;  that  to  almost  every  hydrocar- 
bon there  corresponds  at  least  one  alcohol,  one  aldehyde,  and 
one  acid;  that  the  hydrogen  of  all  these  compounds  can  be 
replaced  in  the  greatest  variety  of  ways  by  chlorine,  bro- 
mine, iodine,  and  by  innumerable  groups  (radicals) ;  and, 
finally,  that  there  are  numerous  other  classes  of  compounds 
which  can  not  even  be  mentioned  here,  he  will  perceive 
that  the  total  number  of  carbon  compounds  must  be  im- 
mense. In  fact,  about  100,000  are  known  at  present,  and 
thousands  of  new  ones  are  being  prepared  every  year.  Most 
of  these  substances  are  entirely  products  of  the  laboratory, 
not  being  found  in  nature,  though  many  others  exist  natu- 
rally in  the  bodies  of  plants  and  animals.  Many  of  them 
have  immense  technical  value  as  medicines,  dyestuffs,  devel- 
opers, and  so  on. 

421.  The  subject  of  the  chemistry  of  the  compounds  of 
carbon  is  usually  called  organic  chemistry.  This  name  is  a 
relic  of  a  time  when  it  was  supposed  that  compounds  like 
those  we  have  been  studying  could  only  be  produced  in  the 
organisms  of  plants  and  animals  under  the  influence  of  the 
life  process.  We  know  now  that  this  was  an  error.  The 
carbon  compounds  can  be  made  in  the  laboratory  like  other 
substances,  but  the  name  "  organic "  chemistry  remains. 
We  can  take  only  a  glance  at  it,  for  it  is  a  vast  and  intricate 
subject,  far  more  extensive  than  the  chemistry  of  all  the 
other  elements  put  together. 


CHAPTER  XLIV 

SOME  ADDITIONAL    CARBON  COMPOUNDS-CHEMICAL 
PROCESSES   OF   THE  ANIMAL   BODY 

422.  The  composition  of  sugar  is  described  by  the  for- 
mula C^H^On.     It  will  be  perceived  that  sugar  is  com- 
posed of  carbon,  hydrogen,  and  oxygen,  and  that  the  last 
two  elements  are  present  in  precisely  the  proportions  in 
which  they  exist  in  water,  two  atoms  of  hydrogen  to  one  of 
oxygen.    A  substance  of  this  kind  is  called  a  carbohydrate. 
The  carbohydrates  all  contain  hydrogen  and  oxygen  in  the 
proportion  in  which  these  elements  exist  in  water,  and  most 
of  them  contain  six  atoms  of  carbon  in  the  molecule,  or  some 
multiple  of  that  number.     They  are  abundantly  contained 
in  plants,  and  make  up  a  most  important  element  in  human 
food. 

423.  Ordinary  sugar,    C^H^O^,   is   also   called   cane- 
sugar  or  saccharose.     It  is  contained  in  the  sap  of  many 
plants,  especially  that  of  the  sugar-cane  and  the  sugar-beet. 
Rock-candy   is    pure    crystallized    saccharose.      Saccharose 
forms  colorless,  transparent  crystals  which  look  like  cubes, 
but  are  really  inclined  prisms.     It  dissolves  in  about  one- 
third  its  weight  of  water  at  ordinary  temperatures,  and  is 
still  more  soluble  in  hot  water.     When  carefully  heated, 
sugar  melts,  and  on  being  allowed  to  cool  solidifies  to  a 
clear,  glassy,  amorphous  mass,  which  is  the  basis  of  many 
kinds  of  candy.    After  a  time  the  mass  crystallizes  and  be- 
comes white  and  opaque.     When  melted  sugar  is  heated 
slightly  above  its  melting-point  it  is  converted  into  a  sub- 


298  ELEMENTARY  CHEMISTRY 

stance  called  caramel.  This  has  a  somewhat  bitter  taste, 
and  is  not  capable  of  crystallization.  It  dissolves  in  watery 
liquids,  producing  an  intense  brown  color,  and  is  largely 
used  for  coloring  liquors. 

424.  Glucose,   or  grape-sugar,   C6H1206,  is  abundantly 
contained  in  grapes  and  many  sweet  fruits,  and  in  honey. 
It  is  largely  manufactured  by  boiling  starch,  C6H1005,  with 
dilute  sulphuric  acid.     The  starch  takes  up  a  molecule  of 
water,  thus-    c ^ ^  +  ^  =  ^^ 

Glucose  is  colorless  and  freely  soluble  in  water.  It  is 
not  so  sweet  as  saccharose,  and  does  not  crystallize  so  read- 
ily. Its  connection  with  the  process  of  fermentation  has 
already  been  referred  to. 

425.  Starch,  C6H1005,  exists  in  plants,  in  round  or  elon- 
gated microscopic  bodies  called  starch-granules,  which  differ 
in  size  and  shape  in  different  plants.    Starch  is  white  and  is 
insoluble  in  cold  water.    When  heated  with  water,  the  gran- 
ules swell  up  and  partially  dissolve,  forming  a  slimy  liquid 
called  starch-paste.    Both  starch-paste  and  solid  starch  are 
colored  blue  by  iodine,  and  this  is  an  extremely  delicate  test 
for  either  substance.     The  molecular  weight  of  starch  can 
not  be  determined  with  our  present  methods.     C6H1005  is 
merely  the  simplest  formula  which  will  describe  its  chemical 
composition.     The  real  formula  is  certainly  a  multiple  of 
this,  probably  a  very  large  multiple.     This  remark  applies 
also  to 

426.  Cellulose,  C6H1005,  which  is  the  substance  of  which 
the  fiber  of  plants  consists.    Wood  and  cotton,  for  instance, 
are  largely  cellulose.    The  best  qualities  of  filter  paper  are 
almost  pure  cellulose.     Cellulose   is   a  white,   amorphous 
mass,  insoluble  in  almost  all  liquids.    It  has  been  found  in 
the  lungs  of  consumptives. 

427.  Ether,  C4H100,  is  the  most  important  anaesthetic. 
It  is  also  called  ethyl  oxide.     The  radical  C2H6  is  called 


SOME  ADDITIONAL  CARBON  COMPOUNDS          299 

ethyl.  Alcohol  is  ethyl  hydroxide,  C2H5OH,  and  ether  bears 
the  same  relation  to  alcohol  that  potassium  oxide  does  to 
potassium  hydroxide — 

Potassium  hydroxide,  KOH,  Ethyl  hydroxide,  C2HBOH, 
Potassium  oxide,  K20.  Ethyl  oxide,  (C2H5)20. 

Ether  is  a  colorless,  inflammable  liquid.  It  is  very  vola- 
tile, and,  when  placed  upon  the  skin,  its  rapid  evaporation 
absorbs  so  much  heat  that  a  sensation  of  intense  cold  is 
felt. 

When  ether  vapor  is  inhaled  it  produces  insensibility. 
When  the  liquid  is  drunk,  a  condition  resembling  intoxica- 
tion, but  more  temporary,  results.  Ether-drinking  is  a  vice 
somewhat  similar  to  alcoholism,  but  still  more  serious  in 
its  results.  It  is  especially  prevalent  in  Ireland. 

428.  Oils  and  fats. — Soap. — There  is  a  whole  series  of 
acids,  beginning  with  formic  and  acetic  acids.  The  higher 
members  are  waxy  solids.  Two  of  these  higher  members 
are  specially  important.  They  are  palmitic  acid,  C16H3202, 
and  stearic  acid,  C18H3602.  We  must  mention  also  oleic 
acid,  C18H3.102,  which  is  a  liquid.  These  three  acids  are 
called  the  fatty  acids,  because,  in  chemical  combination  with 
glycerin,  they  make  up  the  chief  animal  and  vegetable  oils 
and  fats. 

Soaps  are  made  by  boiling  fats  or  oils  with  a  solution 
of  sodium  hydroxide,  NaOH.  The  result  of  this  operation 
is  to  produce,  first,  glycerin,  and  second,  the  sodium  salts  of 
the  three  acids  mentioned  above.  The  latter  make  up  the 
soap.  Soap  is  therefore  composed  of  sodium  palmitate,  so- 
dium stearate,  and  sodium  oleate.  If  more  sodium  hydrox- 
ide was  taken  than  was  needed  to  act  upon  the  fat,  the  soap 
will  also  contain  NaOH.  This  makes  the  soap  unfit  for  use 
upon  the  skin.  The  chapping  and  bleeding  of  the  hands  in 
winter  is  often  due  to  the  use  of  soap  of  this  kind.  The 
presence  of  sodium  hydroxide  can  easily  be  detected  by  pour- 


300  ELEMENTARY  CHEMISTRY 

ing  over  a  freshly  cut  surface  of  the  soap  a  warm  alcoholic 
solution  of  phenol  phthalein.  If  the  surface  becomes  blood- 
red,  the  soap  contains  too  much  sodium  hydroxide  and 
should  not  be  used,  but  if  it  remains  colorless,  or  turns 
only  faint  pink,  it  is  nearly  free  from  it. 

429.  The  albumins  or  proteids. — Albumin  or  proteid  is 
a  name  given  to  a  class  of  substances  all  of  which  contain 
carbon,  hydrogen,  'oxygen,   and  nitrogen.     Most  of  them 
contain  sulphur  and  a  few  contain  iron.    They  are  usually 
insoluble  in  water,  though  soluble  in  solutions  of  certain 
salts.     When  heated  they   decompose,   carbon  being  left, 
and  a  complex  mixture  of  gases  and  vapors  given  off.    This 
mixture  contains  water,  carbon  dioxide,  and  ammonia,  as 
well  as  many  other  substances,  and  possesses  the  unpleas- 
ant odor  of  burned  horn  or  burned  feathers.     The  pro- 
teids are  never-failing  constituents   of   the   organisms   of 
animals   and   plants.     Muscular  tissue,   for   instance,   and 
egg-substance  are  chiefly  composed  of  them. 

Proteids  are  an  absolutely  necessary  constituent  of 
human  food.  It  is  impossible  to  support  life  on  a  diet 
which  does  not  contain  them.  Lean  meat,  eggs,  cheese, 
peas,  and  beans  are  examples  of  foods  rich  in  them.  Life 
can  be  sustained  for  a  long  time  by  food  consisting  chiefly  of 
proteids,  but  such  a  diet  is  very  unwholesome.1 

The  other  two  important  constituents  of  foods  are  fats 
and  carbohydrates.  Cream  and  the  fat  of  meat  are  exam- 
ples of  food  rich  in  fats,  while  rice,  bread,  and  potatoes  con- 
tain large  quantities  of  carbohydrates,  chiefly  starch. 

430.  The    chemical   processes   of   the    body. — Broadly 
speaking,  we  may  say  that  the  proteids  are  the  tissue-forming 
foods,  while  the  fats  and  carbohydrates  are  burned  to  pro- 

1  It  is  a  remarkable  fact  that  some  proteids  are  extremely  poison- 
ous. Thus  the  poisonous  constituents  of  the  venom  of  snakes  consist 
of  substances  that  are  very  closely  similar,  chemically,  to  the  white 
of  egg. 


DIMITRI  IVANOVITCH   MENDELEJEFF 

B.  Siberia,  1834. 
Discoverer  of  the  periodic  law. 


CHEMICAL  PROCESSES  OF  THE  BODY  301 

duce  heat  and  the  energy  of  motion.  The  products  of  this 
burning  are  water  and  carbon  dioxide.  It  takes  place  in 
the  muscles  and  other  solid  tissues  of  the  body — not  in  the 
blood — but  the  arterial  blood  brings  the  oxygen  for  the 
combustion  and  the  venous  blood  carries  away  the  carbon 
dioxide.  We  must  not  conclude  from  this  that  arterial  blood 
contains  no  carbon  dioxide  and  venous  blood  no  oxygen. 
The  brightest  arterial  blood  contains  more  carbon  dioxide 
than  oxygen,  and  the  darkest  venous  blood — even  the  blood 
of  a  suffocated  animal,  which  is  nearly  black — still  contains 
much  oxygen.  But  the  blood  of  the  arteries  contains  half  as 
much  oxygen  again  as  that  of  the  veins,  and  the  venous 
blood  contains  about  one-fifth  more  carbon  dioxide  than  the 
arterial. 

The  oxygen  of  the  blood  is  mainly  carried  by  the  red 
corpuscles,  only  a  trifling  amount  being  dissolved  in  the 
plasma.  The  red  coloring  matter  (haemoglobin)  of  the  cor- 
puscles carries  the  oxygen  in  a  state  of  loose  chemical  com- 
bination. 

This  is  important,  for  if  the  oxygen  was  simply  dis- 
solved in  the  blood,  its  quantity  would  be  directly  propor- 
tional to  the  pressure  of  the  gas  on  the  liquid.  Hence  it 
would  be  impossible  to  climb  high  mountains,  or  to  ascend 
in  balloons  to  elevations  where  the  pressure  of  the  air  is 
small,  for  suffocation  would  result,  and  even  the  natural 
variations  in  pressure  at  the  surface  of  the  earth  might 
cause  grave  disturbances  of  the  oxygen  supply  to  the  tissues. 
On  the  other  hand,  working  in  air  under  pressure — for  in- 
stance, in  caissons,  in  bridge-building,  or  in  diving-bells- 
would  supply  the  organism  too  abundantly  with  oxygen  and 
produce  difficulties  of  another  kind.  As  a  matter  of  fact, 
no  great  reduction  in  the  oxygen  of  the  blood  occurs  until 
the  air-pressure  is  reduced  to  about  one-third  of  its  ordi- 
nary amount.  Then  the  compound  of  oxygen  and  haBmo- 
globin  decomposes  and  symptoms  of  suffocation  appear. 


302  ELEMENTARY  CHEMISTRY 

The  carbon  dioxide  of  the  blood  is  mainly  in  the  plasma, 
only  about  one-third  of  it  being  in  the  red  corpuscles.  Ve- 
nous and  arterial  blood  contain  about  the  same  quantities  of 
nitrogen  and  argon.  Nothing  is  known  about  the  function 
of  these  two  gases  in  the  body. 

The  influence  of  light  on  the  chemical  processes  of 
plants  has  been  discussed.  We  know  that  the  decomposi- 
tion of  carbon  dioxide  in  the  leaves,  the  return  of  oxygen 
to  the  air,  and  the  building  up  of  complex  organic  com- 
pounds, like  starch,  sugar,  and  proteids,  can  not  occur  in 
the  absence  of  sunlight.  Light  has  a  similar  though 
smaller  effect  upon  the  chemical  changes  of  the  animal. 
Experiment  leaves  no  doubt,  for  instance,  that  all  the  com- 
plex chemical  changes  through  which  food  passes  in  the 
human  body,  and  which  result  finally  in  the  production  of 
tissue,  motion,  and  heat,  are  stimulated  by  light,  and  are 
distinctly  retarded  when  the  body  is  placed  in  perfect  dark- 
ness. 

What  is  the  chemical  distinction  between  plants  and 
animals?  This  is  an  interesting  question,  because  the 
chemical  distinction  is  the  only  one  which  will  hold  good 
in  all  cases,  especially  in  considering  microscopic  life.  A 
plant  takes  its  carbon  from  carbon  dioxide,  its  hydrogen 
and  oxygen  mainly  from  water,  its  nitrogen  from  ammonia, 
nitric  acid,  and  various  nitrates,  its  potassium,  sodium, 
phosphorus,  and  so  on,  from  the  soil.  The  raw  material 
which  the  plant  takes  consists  therefore  of  simple  chemical 
compounds,  and  out  of  these  it  builds  up  starch,  proteids, 
and  other  substances  of  great  complexity.  This  kind  of 
chemical  work  is  impossible  for  the  animal.  It  must  have 
its  proteids  and  carbohydrates  supplied  to  it  ready  formed. 
And  its  finished  products  consist  of  carbon  dioxide,  water, 
and  other  comparatively  simple  substances  which  are  again 
fit  to  become  the  food  of  plants. 


CHAPTER   XLV 


THE  PERIODIC  LAW 


431.  Arrangement  of  the  elements  in  series. — Omitting 
hydrogen,  let  us  arrange  the  elements  in  the  order  of  in- 
creasing atomic  weights.  Here  are  the  first  fourteen: 


Lithium 

Beryllium 

Boron 

Carbon 

Nitrogen 

Oxygen 

Fluorine 

Li  =  7 

Be  =  9 

B  =  11 

C  =  12 

N  =  14 

0  =  16 

F  =  19 

Sodium 

Magnesium 

Aluminium 

Silicon 

Phosphorus 

Sulphur 

Chlorine 

Na  =  23 

Mg  =  24 

Al  =  27 

Si  =  28.  5 

P  =  31     |S  =  32 

Cl  =  35.5 

Lithium  is  a  metal  very  similar  to  sodium,  fluorine  is 
a  gas  which  resembles  chlorine  very  closely  and  is  the 
strongest  of  all  the  non-metals.  The  elements  between  the 
two  are  intermediate  in  character:  any  element  is  more 
non-metallic  than  the  element  to  the  left  of  it.  Thus  in 
passing  from  lithium  to  the  right  there  is  a  gradual  loss 
of  metallic  properties,  which  finally,  when  we  arrive  at 
fluorine,  is  complete. 

Now  the  next  element  in  order  is  sodium  (Na  =  23), 
one  of  the  strongest  of  the  metals  chemically.  There  is 
no  gradual  transition  from  fluorine  to  sodium.  We  pass 
at  once  from  the  strongest  non-metal  to  one  of  the  most 
positive  of  the  metals.  There  can  be  no  doubt  that  sodium 
belongs  in  the  same  group  with  lithium,  for  the  two  resem- 
ble each  other  in  a  remarkable  way.  This  is  true  also  of 
the  elements  which  follow  sodium  in  the  second  line;  each 
is  like  the  one  above  it  in  the  first  line.  Magnesium  is  sim- 
ilar to  beryllium,  aluminium  to  boron,  and  so  on.  This 

303 


304  ELEMENTARY  CHEMISTRY 

similarity  is  greater  at  the  ends  than  in  the  middle  of  the 
table.  Sodium  is  more  similar  to  lithium  and  chlorine  to 
fluorine  than  aluminium  is  to  boron  or  silicon  to  carbon. 
Yet  the  similarity  between  these  middle  elements  is  great 
enough  to  show  that  they  belong  together. 

432.  The  "  law  of  octaves." — Thus  these  two  sets  of 
seven  elements  each  exhibit  a  relationship  like  that  of  two 
octaves  in  music: 

1st  octave  C  D  E  F  G  A  B 

2d  octave  c  d  e  f  g  a  b 

In  fact,  when  this  remarkable  arrangement  of  the  elements 
was  first  brought  forward,  it  was  called  the  "  law  of  oc- 
taves "  for  that  reason.  The  properties  of  the  elements 
change  with  increasing  atomic  weight,  and  the  change  is  a 
periodic  one — that  is,  similar  elements  occur  again  and 
again  as  the  atomic  weight  increases,  very  much  as  the 
values  of  the  sine  of  an  angle  repeat  themselves  when  the 
value  of  the  angle  increases. 

This  periodic  change  in  properties  with  increasing 
atomic  weight  is  the  root-idea  of  the  periodic  law,  and,  if 
all  the  elements  behaved  like  the  first  fourteen,  the  whole 
matter  would  be  very  simple.  We  should  arrange  the  ele- 
ments in  the  order  of  increasing  atomic  weights  in  horizon- 
tal lines,  each  containing  seven  elements,  and  those  elements 
falling  in  the  same  vertical  line  would  belong  together 
and  show  similarity  in  properties.  We  shall  see  at  once  that 
the  real  state  of  things  is  more  complicated  than  this. 

433.  Long  and  short  periods. — The  set  of  elements  from 
lithium  to  fluorine  we  call  the  first  short  period,  and  that 
from  sodium  to  chlorine  the  second   short  period.     The 
next  seventeen  elements  in  the  order  of  increasing  atomic 
weights  are  the  following: 

1234567  8 

Potassium  Calcium  Scandium  Titanium  Vanadium  Chromium  Manganese  r- 


K=39  Ca=40  Sc=44  Ti=48  V=51  Cr=52  Mn=55  Iron  Cobalt  Nickel 
Copper  Zinc  Gallium  Germanium  Arsenic  Selenium  Bromine  Pe=56Co=59Ni=58.7 
Cu=63.6  Zn=65.5  Ga=70  Ge=72  As=75  Se=79  Br=80 


THE  PERIODIC  LAW  305 

This  set  begins,  as  we  should  expect  it  to,  with  a  metal 
(potassium)  whose  similarity  to  sodium  and  lithium  is  very 
great.  Farther  along  in  the  first  line  we  discover  that  we 
have  here  a  different  arrangement  from  that  of  the  short 
periods.  Chromium  is  not  very  similar  to  oxygen  and 
sulphur  in  whose  vertical  group  it  falls,  for  it  is  much 
more  metallic  in  character;  while  the  similarity  between 
manganese,  on  the  one  hand,  and  fluorine  and  chlorine  on 
the  other,  is  remote,  manganese  being,  in  most  of  its  chemi- 
cal conduct,  a  metal.  Yet,  though  in  both  cases  the  ele- 
ments differ  from  the  corresponding  ones  of  the  short  peri- 
ods, there  are  still  some  striking  points  of  similarity  which 
justify  us  in  classing  chromium  with  oxygen  and  sulphur, 
and  manganese  with  fluorine  and  chlorine.  One  important 
difference  between  this  arrangement  and  that  of  the  short 
periods  is,  then,  that  at  the  seventh  element— manganese  in 
this  case — the  metallic  properties  have  only  partially  but  by 
no  means  completely  disappeared.  Further,  the  next  metal, 
iron,  is  by  no  means  a  metal  like  sodium,  as  it  should  be 
if  it  stood  at  the  beginning  of  a  new  short  period.  From 
manganese  (Mn  =  55)  to  copper  (Cu  =  63.5)  through  the 
three  elements  of  the  eighth  column — iron,  cobalt,  and 
nickel — there  is  a  gradual  increase  of  metallic  properties 
and  not  a  very  great  increase,  for  copper,  though  a  more 
positive  metal  than  manganese,  bears  no  comparison  to 
potassium  in  that  respect.  In  the  last  seven  elements,  from 
copper  to  bromine  (Br  =  80),  there  is  a  gradual  and  com- 
plete disappearance  of  the  metallic  characteristics.  Bro- 
mine is  an  unmistakable  non-metal,  and  belongs  in  the  same 
group  as  fluorine  and  chlorine. 

This  set  of  seventeen  elements  is  called  the  first  long 
period,  and  the  general  plan  on  which  a  long  period  is  built 
is  this:  First  stands  one  of  the  strongest  positive  metals 
chemically,  and,  following  it,  six  elements  in  which  the 
metallic  qualities  diminish,  but  do  not  completely  disap- 


306  ELEMENTARY  CHEMISTRY 

pear,  so  that  the  seventh  element,  like  manganese,  shows 
mixed  metallic  and  non-metallic  characters.  The  position 
of  the  three  following  elements  is  peculiar.  Their  atomic 
weights  lie  near  together  (compare  the  atomic  weights  of 
iron,  cobalt,  and  nickel),  and  they  resemble  each  other,  but 
we  can  trace  through  them  a  gradual  increase  in  metallic 
properties,  so  that  the  first  member  of  the  second  set  of 
seven — copper,  for  example — is  more  metallic  than  the  last 
member  of  the  first  set  of  seven — manganese,  for  instance. 
Finally,  through  the  remaining  six  elements  the  metallic 
properties  gradually  and  completely  vanish,  so  that  the  last 
member  of  the  long  period  is,  in  all  respects,  a  non-metal. 

434.  Grouping  of  the  elements  according  to  the  periodic 
law. — The  complete  arrangement  of  the  elements  accord- 
ing to  the  periodic  law  is  given  in  the  table.  The  vertical 
columns  are  called  groups,  and  the  student  will  be  pre- 
pared to  find  that  elements  in  the  same  group  resemble 
each  other.  It  is  convenient  to  divide  each  group  into  two 
subgroups,  and  the  resemblance  between  members  of  the 
same  subgroup  is  especially  close.  Thus,  in  the  first  group, 
the  members  of  the  sodium  group  (subgroup  A),  lithium, 
sodium,  potassium,  rubidium,  and  cesium,  resemble  each 
other  far  more  than  they  resemble  the  elements  of  the 
copper  group  (subgroup  B),  copper,  silver,  and  gold.  The 
elements  of  the  first  and  second  groups  are  all  metals,  and 
so  are  all  those  of  the  third  group  except  the  first  one — 
boron.  With  this  exception  the  non-metals  are  all  con- 
tained in  the  fourth,  fifth,  sixth,  and  seventh  groups,  and 
the  strongest  non-metals  stand  at  the  top,  for  in  a  sub- 
group composed  of  non-metals  the  non-metallic  properties 
decrease  with  increase  of  atomic  weight.  This  is  well 
shown  by  the  fact  that  no  non-metal  is  known  having  a 
higher  atomic  weight  than  that  of  iodine  (I  =  127).  In  a 
subgroup  of  metals  the  reverse  is  usually  true — the  higher 
the  atomic  weight  the  more  marked  the  metallic  properties. 


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308  ELEMENTARY  CHEMISTRY 

The  sodium  group  (subgroup  A  of  the  first  group)  con- 
tains those  elements  which  manifest  metallic  properties  in 
the  greatest  perfection,  and  the  most  positive  metal  of 
these,  and,  in  fact,  the  strongest  metal  known,  is  cassium 
(Cs  =  133),  which  has  the  highest  atomic  weight  in  the  sub- 
group. 

435.  Position  of  argon  and  the  allied  elements  in  the 
periodic  system. — Eecently  there  has  been  much  discussion 
as  to  the  place  of  argon,  helium,  neon,  krypton,  and  xenon 
in  the  periodic  system.  All  of  these  elements  exist  in  the 
atmosphere  (p.  100).  They  are  colorless  gases,  and  their 
chief  distinction  is  complete  chemical  inertness — it  is  im- 
possible to  induce  them  to  take  part  in  a  chemical  process, 
and  therefore  we  can  not  classify  them  either  as  metals  or 
non-metals  chemically. 

Now  the  atomic  weight  of  neon  is  about  20.  Therefore 
it  must  follow  fluorine  (F  =  19)  and  precede  sodium 
(Na  =  23)  in  the  table  and  must  fall  in  the  eighth  group. 
But  fluorine  is  the  most  intense  of  non-metals  (p.  251)  and 
sodium  one  of  the  strongest  of  metals.  Neon,  an  element 
of  no  chemical  character  whatever,  comes  in  as  a  natural 
transition  between  these  tivo  extremes. 

Neon  being  thus  provided  with  a  very  satisfactory  posi- 
tion, what  about  the  other  inert  gases  ?  In  the  first  place 
it  is  clear  that  they  must  all  belong  in  the  eighth  group,  if 
neon  is  to  be  placed  there,  for  they  resemble  it  very  closely. 
Argon,  with  its  atomic  weight  of  about  40,  falls  between  the 
strong  non-metal  chlorine  and  the  strong  metal  potassium. 
Krypton  (Kr  =  82)  is  a  transition  between  bromine  (Br  — 
80),  a  distinct  non-metal,  and  rubidium  (Rb  =  85),  a  metal 
of  the  sodium  group.  Finally,  xenon  (X  =  128)  must  be 
placed  between  the  halogen  iodine  (I  =  127)  and  caesium 
(Cs  =  133),  the  most  positive  of  the  sodium  metals.  Thus 
these  elements  find  a  natural  place  in  the  eighth  group 
as  transition  elements  between  the  strongly  non-metallic 


THE  PERIODIC  LAW  309 

halogens  and  the  strongly  metallic  elements  of  the  first 
group. 

436.  Uses  of  the  periodic  law. — The  student  will  notice 
at  once  that  there  is  a  number  of  gaps  in  the  table.    A  gap 
arises  when  we  are  forced  to  leave  a  vacant  space  for  the 
sake  of  preserving  the  arrangement.     Thus,  in  the  second 
long  period,  after  molybdenum    (Mo  =  96)    in  the   sixth 
group,   the   next   known   element   in   order   of   increasing 
atomic  weight  is  ruthenium  (Ru  =  102).     Now  the  whole 
chemical  character  of  this  element  shows  that  it  belongs  in 
the  eighth  group,  and  not  under  manganese.     Further,  if 
we  simply  proceeded  in  order,  placing  ruthenium  under  man- 
ganese, not  only  ruthenium  itself  but  every  following  ele- 
ment would  be  thrown  out  of  place,  and  the  whole  latter 
portion  of  the  table  would  be  disarranged.     Therefore  we 
leave  a  vacant  space  and  preserve  the  arrangement,  believ- 
ing that  the  place  under  manganese  belongs  to  some  un- 
known element  which  has  an  atomic  weight  of  about  100. 
Thirty  years  ago,  when  the  table  was  first  published,  gaps 
were  more  numerous.     The  Russian  chemist  Mendelejeff, 
the  founder  of  the  periodic  law,  gave  it  as  his  opinion  that 
the  vacant  spaces  would  be  filled  by  the  discovery  of  new 
elements,  and  in  several  cases  he  predicted  in  detail  the 
properties  of  these  elements  and  their  principal  compounds 
from  the  place  which  they  ought  to  occupy  in  the  table. 
These  predictions  were  verified,  the  properties  of  the  newly 
discovered  elements  agreeing  with  MendelejefFs  statements 
with  wonderful  closeness — a  striking  proof  that  the  peri- 
odic classification  is  a  real  law  of  nature  and  not  a  mere 
artificial  arrangement. 

437.  The  periodic  law  is  a  great  aid  in  fixing  the  values 
of  atomic  weights. — Perhaps  the  most  important  use  of  the 
periodic  system  is  to  make  us  completely  certain  that  the 
values  of  the  atomic  weights,  as  they  are  accepted  at  pres- 
ent, are  very  close  to  the  correct  ones.     The  place  of  an 


310  ELEMENTARY  CHEMISTRY 

element  in  the  table  fixes  its  atomic  weight  within  narrow 
limits,  and  has  often  given  us  valuable  information  in  this 
direction.  Suppose  a  doubt  to  exist  as  to  whether  the 
atomic  weight  of  sodium  was  23  or  23  X  2  =  46.  Then  the 
fact  that  sodium  is  a  strong  positive  metal  and  resembles 
lithium  very  closely  shows  at  once  that  it  belongs  in  the 
first  group,  and  that  its  atomic  weight  must  lie  between  that 
of  fluorine  (F  =  19)  and  of  magnesium  (Mg  =  24).  This 
leaves  23  as  the  only  possible  value. 

438.  Defects  of  the  periodic  law.— The  classification  is  not 
perfect.  There  is  no  very  satisfactory  place  in  it  for  hydrogen.  Again, 
the  latest  atomic  weight  determinations  indicate  that  in  two  cases 
an  element  has  a  slightly  lower  atomic  weight  than  the  one  which 
precedes  it  in  the  table.     Thus,  nickel  must  follow  cobalt,  but  its 
atomic  weight  is  a  little  smaller,  and  the  same  is  true  of  iodine  and 
tellurium.     These   difficulties  are  serious  and  they  may  result  in 
some  alteration  of  the  present  arrangement,  but  the  fundamental 
fact  that  the  properties  of  the  elements  vary  periodically  with  the 
atomic  weights  is  secure. 

439.  From  the  fact  that  the.  properties  of  the  elements 
vary  with  the  quantity  of  matter  in  the  atom  some  chemists 
have  drawn  the  conclusion  that  the  kind  of  matter  is  the 
same  in  all,  and  that  all  the  elements  are  different  forms  of 
one  universal  substance.     Mendelejeff  himself  is  opposed 
to  this  idea,  but  it  is  difficult  to  deny  that  the  periodic  law 
furnishes  a  strong  excuse  for  it. 

If  the  student  will  carefully  inspect  the  table  he  will 
perceive  that  the  elements  we  have  studied  have  all  been 
arranged  according  to  the  periodic  law.  This,  in  fact,  is  the 
best  possible  classification.  And  the  fact  that  the  elements 
grouped  and  studied  together  were  clearly  similar  to  each 
other  is  positive  proof  of  the  truth  of  the  law  which  fur- 
nished the  classification. 


CHAPTER    XLVI 

THE  HISTORY   OF   CHEMISTRY 

440.  Development  of  chemical  theories. — Without  know- 
ing something  of  the  history  of  our  science  it  is  impos- 
sible to  understand  its  present  state,  for  we  are  likely  to 
fall  into  two  very  serious  errors:  first,  that  of  believing 
that  the  science  originated  all  at  once  in  the  mind  of  the 
writer  of  the  book  we  are  reading;  and,  second,  that  of  think- 
ing that  the  views  which  happen  to  be  held  at  present  upon 
chemical  subjects  are  absolutely  true  and  will  remain  un- 
changed for  all  time.    A  glance  at  the  history  of  chemistry 
will  show  us  that  the  gathering  of  the  vast  wealth  of  facts 
which  the  science  possesses  began  thousands  of  years  ago, 
and  that  the  theories  which  now  aid  us  to  arrange  and  re- 
member the  facts  and  to  discover  new  facts,  have  been  slowly 
developed  from  other  quite  different  theories,  and  will,  in 
all  probability,  be  supplanted,  in  their  turn,  by  others. 

441.  Chemical  knowledge  of  the  ancients. — Among  the 
ancient  Egyptians,  Greeks,  and  Romans  there  was  no  sys- 
tematic chemical  knowledge — nothing  that  could  be  called 
a  science.    Accident,  and  the  needs  of  daily  life,  had  caused 
a  large  number  of  facts  to  become  known,  but  no  attempt 
was  made  to  classify  or  explain  them. 

The  name  of  our  science  is  derived  from  the  Egyptian 
word  Chemi,  which  means  Egypt.  In  that  country  chemis- 
try was  considered  a  sacred  art  and  was  studied  only  by  the 
priesthood.  The  earliest  laboratories  of  which  we  have  any 

311 


312  ELEMENTARY  CHEMISTRY 

record  were  rooms  in  Egyptian  temples  in  which  chemical 
operations  were  secretly  carried  on. 

Gold  and  silver,  which  occur  free  in  nature,  were  prob- 
ably the  first  metals  discovered,  though  copper  was  known 
before  the  dawn  of  history,  and  iron  at  a  very  early  date. 
Lead  was  much  used  by  the  Romans  for  making  water- 
pipes,  and  tin  was  also  known  to  them,  but  they  did  not 
regard  them  as  two  different  metals.  Brass — the  alloy  of 
copper  and  zinc — was  known  in  very  early  times,  but  was 
regarded  simply  as  copper  colored  yellow.  Various  indus- 
tries of  a  chemical  character,  for  instance,  dyeing  and  the 
manufacture  of  glass  and  pottery,  flourished  among  the 
Egyptians,  Romans,  and  Phoenicians.  Soap  was  first  made 
by  the  barbarous  German  tribes  from  wood-ashes  and  ani- 
mal fat.  The  preparation  of  lime  and  of  building-mortar 
from  it  is  very  ancient.  Acetic  acid  (in  vinegar)  was  the 
first  acid  discovered.  Other  compounds  of  carbon  which 
have  been  known  for  more  than  two  thousand  years  are 
sugar,  starch,  and  oil  of  turpentine. 

442.  The  alchemists. — Alchemy  is  the  attempt  to  trans- 
form copper,  mercury,  and  the  other  so-called  "  base  "  met- 
als into  gold  and  silver.  This  effort  began  in  Egypt.  When 
the  Arabs  conquered  Egypt  they  acquired  alchemistic  ideas, 
which  they  carried  with  them  to  Spain.  Thence  alchemy 
penetrated  Europe,  where  it  flourished,  particularly  in  Ger- 
many, for  centuries,  and  thousands  of  men  wasted  their  lives 
in  the  attempt  to  accomplish  the  impossible.  Most  of  them 
were  honest  in  their  belief  that  the  transmutation  of  met- 
als was  possible,  but  others  were  mere  cheats  who  lived  by 
trading  upon  the  credulity  of  princes.  Each  prince  had  his 
court  alchemist,  and — seduced  by  the  hope  of  making  coin- 
age for  his  realm  out  of  alchemistic  gold — expended  vast 
sums  of  money  in  experiments  which  invariably  failed. 

Finally,  when  it  was  discovered  that  the  alchemist  was 
unable  to  accomplish  what  he  had  promised,  he  was  dis- 


THE  HISTORY   OF  CHEMISTRY  313 

missed  in  disgrace,  or,  more  frequently,  treated  with  terri- 
ble severity.  A  favorite  mode  of  punishment  consisted  in 
dressing  the  alchemist  in  a  garment  of  cloth  of  gold  and 
hanging  him  on  a  gilded  gallows. 

The  alchemists  believed  that  the  transformation  of  the 
baser  metals  into  gold  was  to  be  accomplished  by  means 
of  a  substance  called  the  " philosopher's  stone"  The  direc- 
tions which  they  give  for  preparing  this  marvelous  mate- 
rial are  quite  meaningless,  but  with  respect  to  its  properties 
they  speak  more  plainly.  It  was  ruby  red,  transparent, 
crystalline,  and  luminous  in  the  dark.  It  was  very  dense, 
but  brittle  and  easily  powdered,  and  the  powder  glittered 
like  crushed  glass.  Its  possession  conferred  perfect  health, 
and  by  its  proper  internal  use  life  might  be  prolonged  for 
many  centuries.  When  copper  or  some  other  metal  was 
melted  in  a  crucible,  and  some  of  the  "  philosopher's 
stone"  thrown  upon  it,  it  was  completely  transformed 
into  gold. 

The  efforts  of  the  alchemists  to  prepare  this  imaginary 
substance  lasted  for  nearly  fifteen  centuries,  and  every  im- 
aginable material,  not  only  of  mineral,  but  of  vegetable 
and  animal  origin,  was  investigated.  Incidentally  many 
facts  were  discovered  and  methods  devised  which  have  been 
of  great  value  to  our  science.  We  owe  to  the  alchemists  the 
operations  of  filtration,  crystallization,  sublimation,  and 
distillation,  and  the  discovery  of  hydrochloric,  nitric,  and 
sulphuric  acids,  of  arsenic,  antimony,  and  phosphorus,  and 
of  many  important  salts,  like  silver  nitrate  and  potassium 
nitrate.  Most  important  of  all,  we  owe  to  them  the  convic- 
tion that  no  element  can  be  changed  into  any  other;  that 
in  order  to  prepare  gold  or  any  other  element  we  must 
start  with  some  material  Which  contains  it.  It  must  be 
remembered  that  this  law  is  not  self-evident.  It  is  simply 
the  result  of  experience.  To  a  man  who  had  no  chemical 
knowledge  the  transformation  of  copper  into  gold  by  means 


314  ELEMENTARY  CHEMISTRY  • 

of  the  philosopher's  stone  would  not  be  nearly  as  incredible 
as  the  transformation  of  water  into  oxygen  and  hydrogen  by 
the  electric  current. 

It  will  help  us  to  understand  how  the  transmutation 
of  the  metals  could  appear  possible  to  the  alchemists  if  we 
recall  the  fact  that  they  did  not  regard  the  metals  as  ele- 
ments, and,  in  fact,  that  the  word  element  had  not  the  mean- 
ing to  them  that  it  has  to  us.  The  present  definition  of 
the  elements  as  the  constituents  of  compounds,  constitu- 
ents which  can  be  obtained  in  the  free  state,  and  can  not  be 
decomposed  into  simpler  substances,  was  first  given  by  the 
famous  Eobert  Boyle — the  discoverer  of  Boyle's  law — in 
1661.  The  alchemists,  on  the  contrary,  believed  that  the 
metals  were  compounds  of  mercury  and  sulphur.  The  mer- 
cury was  supposed  to  give  them  their  luster  and  the  sulphur 
the  property  of  being  oxidized  or  other  vvdse  altered  when 
heated  in  the  air.  To  any  one  holding  this  belief  the  con- 
version of  one  metal  into  another  appears  a  very  simple 
matter.  Thus  it  was  held  that  tin  and  lead  are  both  com- 
pounds of  mercury  and  sulphur.  The  only  difference  be- 
tween them  is  that  the  tin  contains  more  mercury  and 
less  sulphur  than  the  lead.  Hence  it  was  gravely  asserted 
and  believed  that  lead  could  be  transformed  into  tin  by 
heating  it  with  a  little  mercury  in  a  crucible,  and  no  one 
disproved  the  statement  because  our  present  method  of  in- 
vestigating such  matters  by  experiments  was  not  yet  in 
use.  Gold  and  silver,  on  account  of  their  strong  luster  and 
the  fact  that  they  were  not  affected  by  heat,  were  considered 
to  be  richer  in  mercury  and  poorer  in  sulphur  than  the 
other  metals. 

443.  The  doctrine  of  phlogiston.— Out  of  innumerable 
failures  there  finally  emerged  the  conviction  that  the  object 
of  alchemy  was  impossible,  and  that  alchemistic  labors  were 
a  waste  of  time.  Then  followed  a  period  in  which  it  was 
believed  that  "  the  object  of  chemistry  is  not  to  make  gold 


THE  HISTORY  OF  CHEMISTRY  315 

but  to  prepare  medicines"  (Paracelsus).  During  this  time 
many  new  substances  were  prepared,  and  everything,  both 
new  and  old,  was  tested  to  see  whether  it  could  be  used  in  the 
cure  of  disease.  In  this  way  many  results  of  great  value, 
both  to  chemistry  and  to  medicine,  were  obtained. 

Finally,  about  the  middle  of  the  seventeenth  century, 
chemistry  took  rank  as  a  distinct  science,  with  aims  and 
methods  of  its  own,  and  the  first  problem  to  which  chemists 
turned  their  attention  was  the  explanation  of  the  striking 
phenomena  of  combustion.  What  happens  when  sulphur  or 
a  candle  burns,  or  when  a  metal  is  heated  in  the  air  and 
passes  into  a  lusterless,  earthy  mass  then  called  the  calx? 
It  seemed  clear  that  the  cause  of  combustion  must  be  the 
same  in  all  cases,  and  that  all  combustible  bodies  must  con- 
tain a  common  constituent  which  made  them  combustible. 
To  this  imaginary  substance  the  name  phlogiston  was  given. 
Thus  zinc,  for  example,  is  a  compound  of  the  white  sub- 
stance, zinc  calx — which  we  now  call  zinc  oxide — with  phlo- 
giston. When  the  zinc  is  heated  strongly  the  phlogiston 
escapes  into  the  air,  producing  flame,  and  there  remains  in 
the  vessel  the  other  constituent  of  the  zinc,  a  white  mass 
of  zinc  calx.  So  when  copper  is  heated  the  phlogiston  of 
the  metal  escapes  and  copper  calx,  which  is  black,  remains 
behind.  The  metals  do  not  produce  calxes  when  all  air  is 
excluded  because  the  phlogiston  can  not  escape. 

A  substance  like  charcoal,  which  burns  away  leaving  lit- 
tle or  no  residue,  was  regarded  as  composed  of  pure  or  nearly 
pure  phlogiston.  Thus  it  is  easy  to  see  why  charcoal,  when 
heated  with  the  calxes  (oxides)  of  the  metals,  converts  them 
into  metals.  It  simply  restores  the  phlogiston  which  was 
lost  when  the  metal  was  heated  in  the  first  place.  Thus  tin 
is  a  compound  of  tin  calx  (tin  oxide)  with  phlogiston,  and 
when  tin  is  heated  the  phlogiston  escapes  and  white  luster- 
less  tin  calx  is  left.  But  when  tin  calx  is  mixed  with  char- 
coal— which  is  nearly  pure  phlogiston — and  heated,  the 


316  ELEMENTARY  CHEMISTRY 

charcoal  restores  the  phlogiston  which  had  escaped,  and 
metallic  tin  is  again  obtained. 

For  a  long  time  the  phlogistic  idea  appeared  to  furnish 
a  satisfactory  explanation  of  combustion  phenomena,  but 
during  the  eighteenth  century  it  gradually  became  evident 
that  there  were  grave  defects  in  it.  Thus  there  are  some 
calxes,  like  calx  of  mercury  (mercuric  oxide),  which  when 
heated  alone  are  converted  into  metal,  without  the  aid  of 
charcoal  or  any  other  substance.  It  is  not  clear  whence  the 
phlogiston  comes  in  such  cases.  But  the  great  difficulty 
was  the  fact  that  all  experiments  showed  that  the  calxes 
weighed  more  than  the  metals  from  which  they  were  made. 
Clearly,  if  the  metal  is  a  compound  of  the  calx  with  phlo- 
giston, it  should  weigh  more  than  the  calx. 

444.  The  work  of  Lavoisier. — This  objection  was  a  fatal 
one;  nevertheless  the  theory  of  phlogiston  continued  in 
favor  until  Priestley's  discovery  of  oxygen  supplied  the  key 
to  the  correct  explanation  of  combustion.  Priestley  himself 
never  grasped  the  magnitude  of  his  discovery.  The  expla- 
nation of  the  role  which  oxygen  plays  in  combustion,  in  the 
oxidation  of  metals,  and  in  respiration  was  the  work  of  the 
great  French  chemist  Lavoisier.  It  was  Lavoisier  who 
caused  the  downfall  of  the  theory  of  phlogiston  and  who 
set  up  in  its  place  our  present  views. 

The  law  of  the  indestructibility  of  matter  had  been  ac- 
cepted without  proof  by  many  chemists,  but  we  owe  to  La- 
voisier the  first  experimental  evidence  of  it.  He  sealed  up 
a  quantity  of  tin  in  a  retort,  and,  after  careful  weighing, 
heated  the  tin  for  a  long  time.  When  the  tin  had  combined 
with  the  oxygen  of  the  air  in  the  vessel  he  weighed  again  and 
found  the  weight  unchanged — there  had  been  no  creation 
or  destruction  of  substance  in  the  burning  of  the  tin.  Upon 
opening  the  vessel  it  was  easy  to  show  that  a  portion  of  the 
air  had  disappeared  and  that  the  remainder  was  incapable 
of  supporting  combustion.  This  plan  of  heating  substances 


THE  HISTORY   OF  CHEMISTRY  317 

in  closed  vessels  to  show  that  chemical  changes  produced 
no  alteration  in  weight  was  repeatedly  employed  by  La- 
voisier, and  has  been  much  used  recently  for  the  same 
purpose. 

Lavoisier  was  guillotined  upon  baseless  and  absurd 
charges  during  the  Reign  of  Terror  in  1794. 

445.  The  atomic  theory.  Determination  of  atomic 
weights. — Lavoisier  and  his  predecessors  had  accepted  the 
law  of  definite  proportions — that  the  composition  of  a  com- 
pound is  always  the  same — without  question.  Proust  first 
established  this  law  by  convincing  experiments  in  the  open- 
ing years  of  the  nineteenth  century.  About  the  same  time 
(1804)  occurred  another  event  of  great  importance  to  our 
science — the  statement  of  the  chemical  atomic  theory  by 
John  Dalton. 

The  idea  of  atoms  was  not  new.  It  is  almost  as  old  as 
thought  itself,  and  various  ancient  Greek  philosophers,  es- 
pecially Democritus,  had  framed  theories  of  the  universe 
based  upon  it.  Further,  the  view  that  the  atoms  of  two 
different,  elements  might  unite  to  form  the  particles  of  a 
compound  had  been  stated  by  Boyle  and  employed  by  La- 
voisier. What  Dalton  did  was  to  recognize  for  the  first 
time  that  the  atoms  must  have  definite  relative  weights 
(pp.  94—96),  and  that  by  analyzing  chemical  compounds  it 
is  possible  to  determine  these  weights. 

Chemists  at  once  attacked  the  task  of  determining  the 
atomic  weights.  The  leader  in  this  work  was  the  Swedish 
chemist  Berzelius.  With  tireless  industry  and  wonderful 
skill  he  accomplished  the  immense  labor  of  determining  the 
atomic  weights  of  the  elements  known  at  that  time.  It  is  , 
difficult  for  us  at  present  to  realize  the  conditions  under 
which  Berzelius  and  other  chemists  of  that  time  worked. 
Scarcely  anything  which  Berzelius  needed  could  be  pur- 
chased. Everything  had  to  be  made.  Even  the  three  ordi- 
nary acids  had  to  be  prepared  or  purified  in  the  laboratory. 


318  ELEMENTARY  CHEMISTRY 

There  was  no  gas,  and  the  alcohol  which  Berzelius  needed 
for  his  spirit-lamps  he  was  compelled  to  distill  himself. 
Nevertheless  the  atomic  weights  were  determined  with  an 
accuracy  which,  considering  the  state  of  the  science  at  that 
time,  is  remarkable.  All  of  the  determinations  have  been 
repeated  by  various  chemists,  and  the  atomic  weights  which 
we  now  possess,  particularly  for  the  more  abundant  ele- 
ments, are  quite  exact. 

We  owe  to  Berzelius  also  our  system  of  chemical  sym- 
bols and  formulas,  a  system  which  has  immensely  assisted 
the  progress  of  the  science. 

446.  Throughout  the  first  quarter  of  the  nineteenth  cen- 
tury the  chemical  compounds  of  the  plant  and  animal  king- 
doms were  considered  to  be  totally  different  in  character 
from  other  substances.    They  were  supposed  to  be  produced 
in  the  organism  under  the  influence  of  something  called 
"  vital  force,"  the  result  of  which  was  to  give  rise  to  a 
class  of  substances  distinct  in  character  from  laboratory 
products.     It  was  even  supposed  that  the  simple  laws  of 
definite  and  multiple  proportions  (pp.  33  and  96)  did  not 
apply  to  organic  substances. 

In  1828  Wohler  prepared  artificially  urea,  a  substance 
of  the  composition  of  CON2H4,  found  in  the  urine,  which 
had  hitherto  been  met  with  only  as  a  product  of  the  chemi- 
cal processes  of  the  animal  body.  Since  then  thousands 
of  similar  substances  have  been  made,  and  there  is  no  longer 
any  doubt  that  the  same  laws  hold  good  for  all  chemical 
compounds,  and  that  all  of  the  complicated  products  of  the 
animal  and  plant  organism  will  be  made  artificially.  Many 
*of  them  are  now  manufactured  on  a  large  scale.  Thus  the 
madder-plant  is  no  longer  grown  for  dyeing  purposes,  be- 
cause its  coloring  matter,  alizarin,  can  be  made  much  more 
cheaply  by  chemical  methods.  Artificial  indigo  is  replacing 
the  natural  product. 

447.  The  statement  of  the  periodic  law  by  the  Kussian 


THE  HISTORY  OF  CHEMISTRY  319 

chemist  Mendelejeff  was  made  in  1869.  It  has  been  of  im- 
mense value  to  the  science. 

Among  the  achievements  of  the  closing  years  of  the 
nineteenth  century  may  be  mentioned  the  investigation  of  the 
nature  of  solutions  (pp.  127-131),  the  invention  of  methods 
of  determining  the  molecular  weight  of  dissolved  substances, 
the  work  of  Moissan  with  the  electric  furnace  upon  high 
temperatures  (p.  276),  and  the  discovery  of  argon  and  the 
related  elements — remarkable  on  account  of  their  complete 
chemical  inertness. 

448.  The  practical  and  educational  value  of  chemistry. 
— Chemistry  is  well  worth  studying,  simply  from  the 
standpoint  of  culture,  for  the  sake  of  the  mental  develop- 
ment and  insight  into  the  nature  of  things  which  it  gives. 
But  the  immense  practical  benefits  which  the  science  has 
conferred  and  will  confer  upon  mankind  give  it  claims  of 
another  kind  to  attention.  As  a  result  of  the  work  of  the 
chemist  the  widespread  poisoning  of  the  middle  ages  has 
ceased,  the  productive  capacity  of  our  fields  is  doubled,  the 
ill-smelling  and  useless  coal-tar  has  become  an  inexhaustible 
source  of  delicate  perfumes,  brilliant  dyes,  and  invaluable 
medicines,  surgical  operations  are  performed  without  the 
horrible  suffering  which  once  characterized  them,  and  every 
branch  of  industry  receives  assistance  which  rapidly  becomes 
indispensable  to  its  existence.  What  is  to-day  a  curiosity 
of  the  laboratory  is  to-morrow  manufactured  by  the  ton,  and 
has  become  the  luxury  or  necessity  of  thousands.  No  better 
investment  can  be  made  by  a  country  than  the  proper  equip- 
ment of  chemical  laboratories  for  investigation  and  instruc- 
tion. To  say  that  money  thus  spent  will  return  a  thousand- 
fold is  a  totally  inadequate  statement.  And  to  ask  that 
an  insignificant  fraction  of  the  wealth  which  our  science 'has 
created  shall  be  applied  to  its  further  development  and  dif- 
fusion is  a  very  modest  request. 


APPENDIX 


Abundance  of  the  Elements  in  Nature 


LIST  OF  THE  ELEMENTS  IN 
ORDER  or  ABUNDANCE. 

Composition  of 
the  solid  crust 
of  the  earth. 

Composition  of 
sea-water. 

Composition  of 
the  earth's  crust, 
including  the 
ocean  and  the 
atmosphere. 

Oxygen  .  . 

Per  cent. 
47.29 

Per  cent. 
85.79 

Per  cent. 
49.98 

Silicon 

27  21 

25  30 

Aluminium 

7  81 

7  26 

Iron     .       

5  46 

5  08 

Calcium  

3  77 

0  05 

3  51 

Magnesium  

2  68 

0  14 

2  50 

Sodium  

2  36 

1.14 

2.28 

Potassium 

2  40 

0  04 

2  23 

Hydrogen 

0  20 

10  67 

0  94 

Titanium 

0  33 

0  30 

Carbon  

0  22 

0  21 

Chlorine  

0  01 

2  08 

0  15 

Phosphorus  

0  10 

0  09 

Manganese  

0  08 

0.07 

Sulphur  

0  03 

0.09 

0.04 

Barium 

0  03 

0  03 

Nitrogen    ... 

0  01 

0  02 

Chromium  

0  01 

0  01 

100  per  ct. 

100  per  ct. 

100  per  ct. 

The  crust  of  the  earth  does  not  contain  as  much  as  0.01  per  cent 
of  any  one  of  the  remaining  60  elements.  The  entire  60  make  up  but 
a  small  fraction  of  1  per  cent. 


INDEX 


Acetylene,  288. 
Acids,   126. 

"  Acqua  Toffana,"  223. 
Air,  98. 
Albumin,   300. 
Alchemists,  311,  312,  31° 
Alcohol,  291. 
Aldehyde,  294. 
Alumina,   191. 
Aluminates,  192. 
Aluminium,  189. 

bronze,   191. 

chloride,  194. 

hydroxide,  192. 

oxide,   191. 

sulphate,   192. 
Alums,  192,  193,  194. 
Amalgam,  gold,   160. 

tin,  205. 
Amalgams,  185. 
Amethyst,  198. 
Ammonia,  105,  106,  107. 
Ammonium,  107,  108,  109. 

chloride,  107,  108,  109. 

compounds,   109. 

hydroxide,   110. 

nitrate,  109. 

sulphate,  109. 

Analysis,    use    of   hydrogen    sul- 
phide in,  238. 


Ancients,  chemical  knowledge  of, 

311. 

Anhydride,  sulphuric,  242. 
Animal  body,  chemistry  of,  300, 

301,  302. 
Antimony,  225. 

behavior   of,   in    Marsh's    test, 
226. 

compound    of,   with   hydrogen, 
225,  226. 

trisulphide,  226. 
Aqua  regia,  160. 
Argon,  100. 
Aristotle,  98. 

Arrangement  of  the  elements,  303. 
Arsenic,  220. 

acid,   223. 

allotropy  of,  220. 

disulphide,  223. 

pentasulphide,  224. 

pentoxide,  223. 

trichloride,  77. 

trisulphide,  223. 

white,  223. 
Arsenious  oxide,  223. 
Arsine,  221,  222. 
Asbestos,  193. 
Atmosphere,  98,  99,  100,  101. 

of  sun  and  stars,  34,  35  (foot- 
note). 

821 


322 


ELEMENTARY  CHEMISTRY 


Atom,  93,  94. 
Atomic  theory,  90. 

weight,  94,  119. 

weights,  table  of,  manual, 
Auric  compounds,  161. 

chloride,  162. 

hydroxide,   161. 

oxide,   161. 

Aurous  compounds,  161. 
Avogadro's   hypothesis,    119, 
121,    122,    123. 

Baking-powder,  74. 
Baking-soda,  73. 
Barium,   170. 

peroxide,  47,  48. 

sulphate,  170. 

test  for,  170. 
Bases,  125. 
Basic  salts,  228. 
Berzelius,  317.  . 
Bismuth,  227. ' 

glance,  227. 

nitrate,  228. 

oxynitrate,  228. 

subnitrate,  228. 

sulphide,  227. 
Blast  furnace,  262. 
Bleaching,  79. 

powder,  8$. 
Blende,  177,  178. 
Bluestone,  blue  vitriol,  151. 
Boiling  of  water,  4. 
Bone,  164,  211. 
Borax,   188. 
Boric  acid,  188. 
Boron,  187. 

hydroxide,  188. 

oxide,    188. 
Boyle,  211,  313. 
Brand,  210. 


152. 


120, 


Brass,   146. 

Bromine,  255,  256,  257,  258. 

Bronze,  146. 

Bunsen  burner,  57. 

Cadmium,  179. 

alloys,    179. 

oxide,  180. 

sulphide,  180. 

yellow,   180. 
Caesium,   135. 

hydroxide,  137. 
Calcite   (calcium  carbonate), 
Calcium,  163,  164. 

carbide,  289. 

carbonate,  169. 

carbonate,  dissociation  of, 

chloride,  169. 

dioxide,  166. 

fluoride,   168. 

hydroxide,  "  slaked  lime," 
167. 

iodide,   164. 

oxide,  "  lime,"  165. 

phosphate,  164,  211. 

sulphate,  168. 

sulphide,  167. 
Calomel,    183. 
Calx,  314,  315. 
Carbide   (calcium),  289. 
Carbon,     amorphous,     272, 
274. 

compounds,  series  of,  288, 

crystallized,  diamond,  275. 

crystallized,  graphite,  274. 

dioxide,  278,  279,  280. 

dioxide  in  air,  44,  101. 

dioxide  in  blood,  301,  302. 

dioxide  in  expired  air,  44 

monoxide,  283,  284,  285. 
Carbonado,  275. 


16o. 


166, 


273, 
296 


INDEX 


323 


Carbonates,  280. 

disulphide,  285. 
Carbonic  acid,  280. 
Catalysis,  42,  43,  49,  64,  108,  231. 
Catalytic   action,  42,  43,  49,  64, 

108,  231. 

Caustic    potash     (potassium   hy- 
droxide), 140. 

soda    (sodium  hydroxide),  70. 
Cellulose,  298. 
Cerussite,  209. 
Chalk,  169. 
Chamber  acid,  243. 
Charcoal,  273,  274. 
Chemical  change,  25,  26,  27. 

energy,  280,  281,  282. 

equilibrium,  105. 

Chemistry,    history    of,    Chapter 
XLVI,  311. 

practical  importance  of,  318. 

province  of,  27. 

"Chile    saltpeter"     (sodium    ni- 
trate), 117 
Chloral,  294. 
Chlorates,  87. 
Chloric  acid,  87. 
"Chloride    of    lime"     (bleaching 

powder),  88. 

Chlorides,  preparation  of,  84,  85. 
Chlorine,  75,  76,  77,  78,  79. 

bleaching  action  of,  79. 

compounds  of  oxygen  with,  86. 

heptoxide,  88. 

monoxide,   86. 

peroxide,  86. 
Chromates,  231,  232. 
Chromic  acid,  231. 

chloride,  230. 

oxide,  230. 

salts,  231. 
Chromite,  229. 


Chromium,  229,  230. 

behavior  toward  acids,  230. 

passive  state  of,  230. 

trioxide,  231. 
Cinnabar,  181. 
Clay,  189. 
Coal,  273. 
Cobalt,  268,  269. 

chloride,  269. 

ions,  269. 

Cobaltous  oxide,  268. 
Coke,  273. 

Colloidal  solutions,  224,  266. 
Combination,  chemical,  63. 
Combustion,  Chapter  VII,  51. 
Compounds,  distinction  from  mix- 
tures, Chapter  IV,  31. 
Condensation  of  steam,  2,  3. 
Conductivity,   electrical,  of  met- 
als, 132. 
Copper,  145. 

chloride   (cupric),  150. 

chloride   (cuprous),  148. 

compounds    (cupric),   149,   150, 
151. 

compounds   (cuprous),  148. 

oxide  (cupric),  147. 

oxide    (cuprous),   147. 

sulphate    (cupric),   151. 

sulphide   (cuprous),  145. 
Copperas,  265. 
Corrosive     sublimate      (mercuric 

chloride),  184. 
Corundum,    191. 
Cryolite,  189,  190. 
Crystallization,  13,  14. 

of    carbon    forming    diamond, 

276,  277. 
Cupellation,  152. 
Cupric,  bromide,  150. 
!      chloride,  150. 


824 


ELEMENTARY  CHEMISTRY 


Cupric  iodide,  151. 

ions,  149. 

oxide,  147,  148. 

sulphate,  151. 
Cuprous  bromide,  148. 

chloride,  148. 

iodide,  148. 

ions,  149. 

oxide,  147. 
Cyanogen,  124. 

Daguerreotype  process,  155. 
Dalton,  94,  316. 
Davy,  68. 
Decomposition,  63. 
Diamond,  275,  276,  277. 
Disinfection,  291. 
Dissociation,  10,  114. 

of  calcium  carbonate,  165. 

electrolytic,      Chapter      XVII, 

125. 
Distillation,  2,  3. 

of  alcohol,  293. 
Dolomite,  172. 

Earth's     crust,     composition     of, 

Appendix   (table). 
Electrical  conductivity,  128,  130, 

131,  132. 

Electric  furnace,  276. 
Electrolysis,  128,  129,  131. 

of  hydrochloric  acid,  80. 

of  water,  6. 
Elements,  33,  34. 
Emery,  191. 
Epsom  salts,  175. 
Equations,  61. 
Equilibrium,  chemical,  105. 
Ethane,  288. 
Ether,    298. 
Ethyl  alcohol,  291,  292. 


Explosive    mixtures,    287     (foot- 
note). 

Feldspar,  189. 
Fermentation,  291. 
Ferric  chloride,  267. 

hydroxide,  266. 

oxide,  266. 

salts,  266. 
Ferrous  oxide,  264. 

salts,  265. 

sulphate,  265. 

sulphide,  265. 
Fire-damp,  287. 
Flame,  53,  54. 

color,  53. 

reversal  of,  54,  55. 

structure  of,  56. 
Flash-light,  173,   174,  191. 
Flowers  of  sulphur,  235. 
Fluorine,  251,  252,  253. 
Fluor  spar,  251. 
Formaldehyde,  291. 
Formula,  59,  60. 

structural,  142. 
Freezing-point   of   water,  4. 
Fuels,  273. 
Furnace,   electric,   276. 

Galenite,  208. 
Galvanized  iron,  178. 
German  silver,  269. 
Glass,   199. 

etching  on,  254. 

Glauber's      salt       (sodium      sul- 
phate), 14. 
Glucose,  298. 
Gold,  159,  160,  161. 

compounds.       See    Auric    and 
Aurous  Compounds. 

plating,    16C 


INDEX 


325 


Graphite,  274. 
Gunpowder,  117. 
Gypsum,  168. 

Haematite,  266. 
Helium,   100. 
Horn  silver,  152. 
Hydraulic  mining,  159. 
Hydrazine,  110. 
Hydrazoic  acid,  110. 
Hydriodic  acid,  259. 
Hydrobromic  acid,  257. 
Hydrocarbons,  286,  287,  288,  289. 

substitution    products    of,    289, 

290. 
Hydrochloric  acid,  78,  79,  80,  81, 

82,  83. 

Hydrofluoric  acid,  253. 
Hydrogen,  Chapter  V,  35. 

antimonide   (stibine),  225,  226. 

arsenide  (arsine),  221,  222. 

peroxide,  48,  49,  50. 

persulphide,  239. 

phosphide  (phosphine),214,  215. 

silicide,  196,  197. 
Hydrolysis,  127. 
Hydroxyl,  110,  130. 
"Hypo"    (sodium   thiosulphate), 

157,  245. 
Hypochlorous  acid,  134. 

Illuminating  gas,  273. 

ammonia  from,  106. 
Indicators,  125. 
lodic  acid,  260. 
Iodine,  258,  259,  260. 

test  for,  44,  45,  260. 
Iodine  pentoxide,  260. 
lonisation,  128. 
Ions,  128. 

of  acids,  129. 


Ions  of  bases,  130. 

of  salts,  130. 
Iron,  260. 

compounds.    See  ferric  and  fer- 
rous compounds. 
Isomorphism,  193. 

Krypton,  100. 
place  of,  in  periodic  table,  308. 

Lamp-black,   274. 
Lavoisier,  316. 

Law  of  Avogadro,  Chapter  XVI, 
119. 

definite  proportions,  33,  95. 

multiple  proportions,  96. 

octaves,  304. 
Lead,  206. 

carbonate,  209. 

chambers,    for    sulphuric    acid 
process,  242. 

chloride,  209. 

chromate       (chrome      yellow), 
231. 

dioxide,  208. 

glass,  199. 

nitrate,  208. 

oxide   (litharge),  207. 

pipes,  207. 

poisoning,  207. 

sulphate,  209. 

sulphide,  208. 

white,  209. 
Lime,  165. 

slaked,  166. 
Limestone,  169. 
Litharge,  207. 
Lithium,  137. 

Magnesia,   174. 
Magnalium,  173. 
Magnesium,  172. 


ELEMENTARY  CHEMISTRY 


Magnesium  carbonate,  172. 

chloride,   176. 

oxide,  174. 

sulphate,  174. 

sulphate    crystallized     (Epsom 
salts),   175. 

sulphate,  hydrates  of,  175. 
Magnetite,  267. 
Malleable  iron,  263. 
Manganese,  246,  247. 

dioxide,  246. 

heptoxide,  248. 

ions,  247. 

Manganic  salts,  248. 
Manganous  oxide,  247. 

sulphate,  247. 
Marble,  169. 
Marsh  gas,  286. 
Marsh's   test,   221,   222. 
Massicot,  208. 
Matches,  214. 
Matter,  changes  in,  23. 

atomic  constitution  of,  Chapter 

XII,  90. 

Mendelejeff,  309. 

Mendelejeff's      periodic      system, 
Chapter  XLV,  303. 

table,  306. 
Mercuric  chloride,  184. 

iodide,  24,  185. 

ions,  182,  183. 

oxide,  25,  26,  183. 

sulphide,  184. 
Mercurous  chloride,  183. 

ions,  182,  183. 

nitrate,  183. 
Mercury,  181,  182. 

alloys  of    (amalgams),   185. 
Metallurgy  of  iron,  262,  263,  264. 
Metaphosphoric  acid,  216. 
Methane,  286. 


Methyl  alcohol,  290. 

Mixture,     distinction     of,     from 

compound,  Chapter  IV,  31. 
Moissan,  276,  319. 
Molecular  weight,   119,  120,  123, 

124. 

Molecules,  90,  91,  92. 
Mono-sodium  carbonate,  73. 
Mortar,  167. 

Nascent  state,  221    (foot-note). 
Negative      (photography),      156, 

157. 
Neon,  100. 

place    of,    in    periodic    system, 

308. 
Nickel,  268,  269. 

plating,   269. 

salts,  269. 

Nickelous  oxide,  269. 
Nitric  acid,  115. 

oxide,  112. 
Nitrogen,  Chapter  XIII,  98. 

bromide,  118. 

chloride,  118. 

group,  210. 

halogen   compounds  of,   118. 

hydrogen  compounds  of,  Chap- 
ter XIV,  105. 

iodide,  118. 

oxygen  compounds  of,  Chapter 
XV,  111. 

pentoxide,  114. 

peroxide,  113. 

trioxide,  111. 
Nitrous  acid,  117. 

oxide,  111. 

Osmium,  270. 

Oxidation   of   phosphorus,   212. 
of  ferrous  compounds,  264,  265. 


INDEX 


327 


Oxidation  by  ozone,  48. 
Oxide,  43,  58. 
"  Oxidized  "  silver,  154. 
Oxygen,  Chapter  VI,  40. 

group,  234. 
Oxy haemoglobin,  301. 
Oxyhydrogen  blowpipe,  38,  39. 
Ozone,  44,  45,  46,  47,  48. 

Palladium,  270. 

Partial   decomposition.     See  Dis- 
sociation. 

Passivity  of  chromium,  230. 
Peat,  273. 
Periodic    system,    Chapiter    XLV, 

303. 

Permanganic  acid,  248. 
Petroleum,  288. 
Phlogiston,  314. 
Phosphine,  214,  215. 
Phosphonium  iodide,  215. 
Phosphorous  acid,  217. 
Phosphorus,  210. 
Pig  iron,  263. 

acids  of,  210. 

chlorides  of,  219. 
Plaster  of  Paris,  168. 
Platinum,  270. 

black,  270. 

dichloride,  271. 

dioxide,  271. 

monoxide,  271. 

tetrachloride,   271. 

oxides  of,  216. 

oxy chloride,  219. 

pentachloride,  219. 

pentasulphide,  218. 

trichloride,  219. 
Photography,  155. 
Physical    and    chemical    change, 
Chapter  III,  23. 


Platinum  metals,  270. 
Polymer,  114. 
Potassium,  138. 

carbonate,    139. 

chlorate,  41,  87. 

chloride,   143. 

cyanide,  127. 

dichromate,  232. 

hydroxide,  140. 

nitrate,  116. 

oxide,  140. 

permanganate,  248. 

sulphides,  141,  142. 

tetroxide,  140. 
Priestley,  316. 
Primordial  matter,  310. 
Proteid,  300. 
Puddling,  263. 
Pyrolusite,  246. 

Quartz,  197,  198. 
Quicksilver  (mercury),  181. 

Radio-activity,  232. 

Realgar,  223. 

Rhodium,  270. 

Rock  crystal  (quartz),  197,  198. 

Rock  salt,  66. 

Ruby,  191. 

Rust,  262. 

Saccharose,  297. 

Sal-ammoniac    (ammonium   chlo- 
ride), 108,  109. 
Salt,  table,  64,  65,  66,  67. 
Saltpeter.      See    Potassium    Ni- 
trate. 
Salts,  126. 

basic,  228. 

ions  of,  130,  131. 

solutions  of,  130,  131. 


328 


ELEMENTARY  CHEMISTRY 


Sapphire,  191. 
Scheele,  76. 
Selenic  acid,  245. 
Selenium,  245. 
Sensitive  film,  156,  158. 
Serpentine,  2,   195. 
Silica,  197. 
Silicates,  199. 
Silicic  acid,  198. 
Silicon,  195. 

fluoride,  254. 

oxide  (quartz),  197. 
Silver,  152,  153. 

bromide,  155. 

chloride,   155. 

dioxide,   154. 

fluoride,  154. 

iodide,  155. 

nitrate,  154. 
Smalt,  269. 
Soap,  299. 

testing  of,  299. 

Soda,  baking.     See  Mono-sodium 
Carbonate. 

caustic.     See  Sodium  Hydrox- 
ide. 

washing.     See  Sodium  Carbon- 
ate. 

Soda-glass,  199. 
Sodium,  68. 

borate.    See  Borax. 

carbonate,  71,  72. 

chloride,  64. 

hydroxide,  70. 

(mono-)  carbonate,  72,  74. 

nitrate,  117. 

nitrite,  117. 

oxide,  69. 

peroxide,  69. 

silicate  (water-glass),  200. 

sulphate,  14. 


Sodium  sulphite,  241. 

thiosulphate     ("hypo"),     157, 

245. 
Solubility,  12. 

effect  of  temperature  on,  13. 
Solution,  Chapter  II,  12. 

supersaturated,  14. 

theory   of,   127,   128,    129,    130, 

131. 

Soot,  274. 

Stannic   acid.     See   Stannic  Hy- 
droxide. 

chloride,  205. 

hydroxide,  203,  204. 

oxide,  203. 

sulphide,  204. 
Stannous  chloride,  203. 
Starch^  298. 
Stas,  153. 
Steel,  264. 
Stibine,  225,  226. 
Stibnite,  225. 
Strength  of  acids,  129. 

of  bases,  130. 
Strong    acids     and    bases,     129, 

130. 
Strontium,  170. 

nitrate,  170. 
Sugar,  297. 
Sulphites,  241. 
Sulphur,  234. 
Sulphur  dioxide,  239. 
Sulphuretted    hydrogen     (hydro- 
gen sulphide),  237,  238. 
Sulphuric  acid,  242,  243,  244. 

anhydride,  242. 
Sulphurous  acid,  240,  241. 
Sulphur  trioxide,  241. 

Tellurium,  245. 
Thiosulphuric  acid,  246. 


INDEX 


329 


Tin,  Chapter  XXVIII,  201. 

compounds.      See    Stannic    and 
Stannous  Compounds. 

foil,  203. 
Type  metal,  225. 

Uranium,  232. 

Valence,  137,  138. 

Vapor    pressure    of    water.      See 

Table  in  Manual,  153. 
Velocity  of  chemical  changes,  ef- 
fect of  temperature  on,  51. 
Vermilion,   184. 
Vinegar,  294. 
Vitriol,  blue,  151. 
green,  265. 

oil    of    (sulphuric    acid),    242, 
243,  244. 


Water,  Chapter  1, 1. 

glass,  200. 
Wines,  percentage  of  alcohol  in, 

292. 
Wrought  iron,  263. 

Xenon,  100,  308. 
Yeast,  292. 

Zinc,  177. 

blende,  178. 

carbonate,  177. 

oxide,  178. 

sulphate,  179. 

sulphide,  178. 

white,  178. 
Zymase,  292. 


(5) 


THE  END 


ELEMENTARY 
CHEMISTRY 

PART  II 

EXPERIMENTAL  WORK 


BY 

ROBERT  HART  BRADBURY,  A.M.,  PH.D. 

TEACHER   OF   CHEMISTRY,  CENTRAL   MANUAL  TRAINING-SCHOOL, 

PHILADELPHIA  ;     FORMERLY     LECTURER    ON    PHYSICAL 

CHEMISTRY,     DEPARTMENT     OF     PHILOSOPHY, 

UNIVERSITY     OF    PENNSYLVANIA 


NEW    YORK 

D.     APPLETON     AND    COMPANY 
1909 


COPYRIGHT,  1903 
BY  D.    APPLETON   AND   COMPANY 


PREFACE 


IN  the  preparation  of  the  Laboratory  Manual  I  have 
tried  to  make  the  directions  as  complete  and  clear  as  pos- 
sible. Wherever  there  is  any  danger  in  carrying  out  an  ex- 
periment, the  necessary  precautions  are  explicitly  stated, 
even  at  the  cost  of  some  repetition. 

Most  of  the  experiments — but  not  all  of  them — are 
intended  for  individual  laboratory  work.  This  is  a  matter 
which  will  vary  with  the  equipment  of  the  school  and  the 
size  of  the  classes.  The  following  is  a  list  of  the  numbers 
of  the  experiments  which  I  have  found  to  be  better  fitted 
for  lecture-table  or  laboratory  demonstrations  than  for 
individual  work,  but  it  is  merely  the  result  of  my  own  ex- 
perience, and  is  not  intended  to  have  any  prescriptive 
implication : 

Numbers  7,  9, 10,  27,  31, 40,  49a,  50,  51,  60, 128, 129, 133. 

The  questions  placed  after  each  chapter  are  based 
partly  on  the  laboratory  exercise  and  partly  on  the  corre- 
sponding chapter  of  Part  I.  The  student  should  not  fall 
into  the  error  of  supposing  that  his  own  experimental 
work  supplies  all  the  necessary  data  for  answering  them. 

R.  H.  B. 

CENTRAL  MANUAL  TRAINING  SCHOOL,  PHILA. 


CONTENTS* 


CHAPTER  PAOI 

INTRODUCTION — EXPERIMENTAL  WORK  ....  1 

I.— WATER 5 

II. — SOLUTION        .........  16 

III. — PHYSICAL  AND  CHEMICAL  CHANGE — LAW  OP  THE  IN- 
DESTRUCTIBILITY OF  MATTER 21 

IV. — MIXTURE — ELEMENT — COMPOUND— THE  LAW  OF  DEFI- 
NITE PROPORTIONS       , 25 

V.— HYDROGEN       .........  28 

VI. — OXYGEN  AND  HYDROGEN  PEROXIDE   ....  33 

VII. — COMBUSTION        ,      44 

VIII.— PROBLEMS 46 

IX. — SALT  AND  SODIUM 47 

X. — CHLORINE 52 

XI. — CHLORIDES  —  COMPOUNDS    OF    CHLORINE    CONTAINING 

OXYGEN 61 

XII. — QUESTIONS 64 

XIII. — THE  ATMOSPHERE — NlTROGEN           .....  65 

XIV.— AMMONIA 72 

XV. — COMPOUNDS  OF  NITROGEN  AND  OXYGEN.       ...  76 

XVI. — ATOMIC  AND  MOLECULAR  WEIGHTS — AVOGADRO'S  LAW  .  82 

XVII. — ACIDS,  BASES,  AND   SALTS — METALS   AND  NON-METALS  .  84 

XVIII.— THE  SODIUM  GROUP 90 

XIX.— THE  COPPER  GROUP 92 

XX.— SILVER ....  94 

*  Chapters  correspond  to  those  in  Part  I. 

vii 


viii  ELEMENTARY  CHEMISTRY 

CHAPTER  PAGE 

XXL— PROBLEMS 96 

XXII. — THE  CALCIUM  GROUP 97 

XXIIL— MAGNESIUM 98 

XXIV. — ZlNC   AND   CADMIUM 100 

XXV.— MERCURY     .                        102 

XXVI. — BORON  AND  ALUMINIUM 104 

XXVIL— SILICON 107 

XXVIIL— TIN 108 

XXIX.— LEAD 110 

XXX  AND  XXXI. — QUESTIONS  AND  PROBLEMS      .        .        .111 
XXXII  AND  XXXIII.— ARSENIC  AND  ANTIMONY       .        .        .112 

XXXIV.— CHROMIUM 114 

XXXV.— SULPHUR 115 

XXXVI.— NO   EXPERIMENTS 122 

XXXVII.-FLUORINE 122 

XXXVIII.— BROMINE  AND  IODINE 123 

XXXIX.— IRON 127 

XL.— PLATINUM 130 

XLL— CARBON 131 

XLII. — CARBON  DIOXIDE  AND  CARBON  MONOXIDE  .        .        .132 

XLIII. — SOME  CARBON  COMPOUNDS 136 

XLIV. — ADDITIONAL  CARBON  COMPOUNDS       ....  139 

XLV. — NO  EXPERIMENTS 140 

XLVI.— NO   EXPERIMENTS 141 

APPENDIX — CALCULATION  OF  THE  EFFECT  OF  TEM- 
PERATURE, PRESSURE,  AND  WATER-VAPOR   ON   THE 

VOLUMES  OF  GASES 145 

LOGARITHMS  156 


ELEMENTAKY  CHEMISTEY 


PAET  II 

EXPERIMENTAL  WORK 

General  suggestions. — Chemical  laboratory  work  is  by 
no  means  free  from  danger.  The  eyes,  especially,  are 
likely  to  be  injured  if  the  work  is  not  done  with  proper 
care.  The  secret  of  safety  is  the  accurate  observance  of 
the  directions  given  for  the  performance  of  the  experi- 
ment. Before  carrying  out  an  experiment  the  student 
should  carefully  read  the  directions  and  should  observe 
them  scrupulously,  asking  for  information  on  any  points 
which  are  not  clear. 

Neatness  is  essential  to  success.  Experiments  made  in 
dirty  test-tubes,  beakers,  or  mortars,  are  worthless,  mis- 
leading, and  often  dangerous.  Every  piece  of  apparatus 
employed  must  be  spotlessly  clean  and  must  be  cleaned 
carefully  after  the  work  is  done. 

The  sinks  in  the  laboratory  tables  are  intended  only 
for  liquids.  Solids,  like  bits  of  broken  glass,  match-sticks, 
and  paper,  must  never  be  thrown  into  them. 

The  student  must  be  continually  on  his  guard  against 
the  tendency — almost  universal  with  beginners — to  take 
too  much  of  the  various  chemicals  required  in  the  experi- 
ments. One  objection  to  this  is  waste  of  material.  An- 
other is  waste  of  time,  for  it  always  requires  a  much  longer 
time  to  carry  out  a  process  with  a  large  quantity  of  sub- 
stance than  with  a  small  quantity.  Finally,  there  are  many 


2  ELEMENTARY  CHEMISTRY 

experiments  which  are  perfectly  safe  on  the  small  scale 
but  become  highly  dangerous  when  the  quantity  is  in- 
creased. As  a  rule,  the  quantity  of  material  to  be  taken 
is  indicated  in  the  directions,  and  nothing  but  loss  of  time 
and  danger  result  from  taking  more.  Where  no  amount 
is  stated,  take  the  smallest  that  you  can  conveniently  work 
with. 

Anything  which  is  spattered  into  the  eyes  must  be  re- 
moved instantly  by  copious  washing  with  water.  Acid 
splashed  upon  the  skin  should  be  washed  off  at  once  with 
abundant  water,  and  if  a  wound  has  resulted  from  its  ac- 
tion, it  should  be  treated  with  a  paste  made  of  sodium 
acid  carbonate  (baking  soda)  and  water.  The  same  paste 
may  be  applied  to  burns. 

PRELIMINARY  EXERCISE. — THE  BUNSEN  BURNER;  GLASS 

ROD  AND  TUBING 

A.  The  Bunsen  burner. — Examine  the  burner.  Close 
the  holes  at  the  base  and  light  it.  Hold  a  piece  of  glass 
tubing  in  the  flame  for  a  time.  Open  the 
holes  and  describe  the  change  which  the  flame 
undergoes.  Hold  a  piece  of  glass  tubing  in 
this  flame.  Which  flame  is  hotter  ?  Which  is 
cleaner?  Which  is  least  affected  by  drafts? 
What  is  the  cause  of  the  difference  between 
the  two  flames?  Which  is  the  hottest  part 
of  the  blue  flame? 

Is  the  blue  flame  hollow  or  solid?  Ob- 
tain facts  to  answer  this  question  by  thrust- 
jng  a  match-stick  horizontally  through  the 
flame  near  the  burner.  Support  a  match  by  means  of  a 
pin,  as  shown  in  Fig.  1,  and  light  the  burner.  Remove 
the  match,  and  hold  in  the  center  of  the  blue  flame  near  the 
burner  one  end  of  a  glass  tube  open  at  both  ends.  The 
tube  should  be  inclined  obliquely  upward.  Hold  a  lighted 


EXPERIMENTAL  WORK  3 

match  to  the  upper  end  of  the  tube.  What  does  this  show 
regarding  the  nature  of  the  interior  of  the  flame? 

Open  the  holes  at  the  base  as  wide  as  possible,  and 
gradually  turn  off  the  gas  until  the  flame  strikes  back. 
Why  does  the  flame  only  strike  back  when  the  current 
of  gas  up  the  chimney  becomes  slow?  While  the  flame 
is  still  burning  below,  turn  on  the  gas  and  light  it  above. 
Is  this  flame  suitable  for  use?  How  can  the  burner  be 
made  to  give  the  blue  flame  again? 

The  flame  must  never  be  allowed  to  burn  below,  since 
it  gives  off  poisonous  gases.  Remember  in  using  the 
burner  that — except  where  high  temperatures  are  required 
— a  small  flame  is  better  than  a  large  one.  Do  not  light 
the  burner  until  you  are  ready  to  use  it,  and  always  turn 
it  down  or  extinguish  it  when  it  is  not  in  use. 

Take  the  burner  apart,  make  drawings  of  the  parts, 
and  explain  the  function  of  each. 


PIG.  2. 

B.  Glass  rod. — With  a  triangular  file  make  a  notch 
on  a  piece  of  glass  rod  15  centimeters  1  (6  inches)  from 
the  end.  One  sharp  stroke  of  the  file  is  sufficient.  Hold- 
ing the  rod  as  indicated  in  Fig.  2  endeavor  to  bend  it 
away  from  the  notch,  and  it  will  break  off  at  that  point. 

1  Hereafter  the  abbreviation  cm.  for  centimeters  will  be  used. 
23 


4  ELEMENTARY  CHEMISTRY 

Cut  off  three  such  pieces.  Since  the  ends  of  the  rods  are 
jagged  and  inconvenient  to  handle,  round  both  ends  of 
each  rod  by  holding  it  in  the  Bunsen  flame  and  rotating 
the  rod  gently.  Support  the  rods  by  the  middle  on  your 
test-tube  rack  until  they  cool.  (Hot  glassware  or  hot 
apparatus  of  any  kind  must  never  be  laid  on  the  desk 
or  put  away  under  it.) 

C.  Glass  tubing. — Examine  a  piece  of  hard  and  a  piece 
of  soft  glass  tubing.  Carefully  note  and  record  the  dif- 
ferences in  appearance  between  them.  Select  for  yourself 
a  piece  of  each  kind  of  glass  from  the  main  stock.  Cut, 
just  as  you  cut  the  glass  rod,  two  pieces  of  the  soft  glass, 
each  20  cm.  (8  inches)  in  length.  Cut  one  piece  of  hard 
glass  tubing  of  the  same  length. 

Hold  the  middle  of  your  piece  of  hard  glass  tubing  in 
the  flame,  turning  it  gently,  and  when  it  becomes  red- 
hot,  gently  and  slowly  draw  the  two  portions  asunder. 
Do  not  twist  the  tubes.  The  pull  must  be  straight.  Let 
the  two  tubes  cool  and  use  them  to  study  the  effect  of 
heat  upon  a  fragment  of  wood  (match-stick)  and  a  little 
sugar. 

In  the  same  way  make  four  sealed  tubes  of  soft  glass. 
The  temperature  required  is  not  so  high,  and  the  glass 
must  be  removed  from  the  flame  before  drawing  it 
out.  Then  the  thin  middle  portion  must  be  returned 
to  the  flame  and  melted  and  the  two  tubes  formed  drawn 
apart. 

Use  these  tubes  for  studying  the  effect  of  heat  upon 
a  little  paper,  a  crystal  of  iodine,  a  fragment  of  sulphur, 
and  a  small  piece  of  starch.  Never  heat  anything  in  a 
tube  sealed  at  BOTH  ends,  since  this  would  cause  explo- 
sions. Soft  glass  tubing  is  used  for  all  ordinary  purposes ; 
and  hard  glass  tubing — which  is  much  more  expensive 
and  more  difficult  to  manipulate — only  when  high  tem- 
peratures are  to  be  applied. 


EXPERIMENTAL  WORK  5 

Bending  glass  tubing. — For  bending  use  the  flame  of 
a  wing-top  burner,  never  the  Bunsen  flame.  Hold  the 
tube — which  should  be  of  soft  glass — so  that  the  flame 
heats  as  long  a  portion  as  possible,  and  rotate  it  so  that 
it  is  evenly  heated.  When  sufficiently  hot  remove  it  from 
the  flame  and  make  the  bend.  When  the  tube  is  perfectly 
cold  remove  the  soot  by  wiping  the  outside  with  paper. 

In  this  way,  bend  a  soft  glass  tube  into  an  acute  angle 
(Fig.  3).  Select  a  piece  of  soft  glass  tubing  about  50 
cm.  (20  inches)  long  and  bend  it  twice  at  right  angles 


FIG.  3. 


PIG.  4. 


(Fig.  4).  This  must  be  done  so  that  when  the  double 
bend  is  laid  upon  the  table  every  part  of  it  will  touch  the 
surface  of  the  latter,  or,  in  other  words,  the  two  limbs 
must  be  in  the  same  plane.  Round  the  sharp  ends  of  the 
bent  tube  by  holding  them  a  short  time  in  the  flame — not 
long  enough  to  cause  them  to  collapse. 


CHAPTER   I 

WATER 

EXPERIMENT  1. — Place  a  small  piece  of  potato  in  a 
dry  test-tube.  Clamp  the  tube  horizontally  and  heat  gently 
with  a  small  flame.  As  soon  as  a  positive  result  is  ob- 
tained stop  the  experiment. 


6  ELEMENTARY  CHEMISTRY 

What  is  the  result?  Other  vegetable  substances  would 
behave  in  the  same  way.  Draw  a  conclusion  regarding 
the  existence  of  water  in  living  things.  Mention  some 
vegetable  products  iiT*which  the  presence  of  abundant 
water  is  evident  on  mere  inspection. 

EXPERIMENT  2. — Crush  some  ice  in  a  mortar  and  half 
fill  a  small  round-bottomed  flask  with  it.  Dry  the  out- 
side of  the  flask  carefully  with  a  towel  and  clamp  it  about 
25  cm.  (10  inches)  above  the  desk.  Directly  under  it 
place  a  clean,  dry  beaker.  Let  the  apparatus  stand  for  an 
hour  or  mo-re  while  you  go  on  with  other  work.  What 
does  the  result  prove?  How  does  warm  air  differ  from 
cold  air  in  its  capacity  for  water  ?  What  is  dew  and  how 
is  it  formed?  What  is  humidity  and  how  does  it  affect 
our  sensations  in  warm  weather?  Suppose  the  flask  had 
contained  a  mixture  of  ice  and  salt,  what  difference  would 
have  been  noticed  in  the  result  ?  If  water  should  suddenly 
cease  to  evaporate  into  the  air,  what  changes  would  occur 
upon  the  earth's  surface?  Sketch  the  apparatus  in  your 
note-book. 

EXPERIMENT  3. — Carefully  evaporate  a  little  faucet 
water  to  dryness  in  a  perfectly  clean  porcelain  dish.  Is 
the  water  pure?  If  not,  what  is  the  source  of  the  im- 
purities? Evaporate  a  little  distilled  water  in  the  same 
way.  What  is  the  difference? 

EXPERIMENT  4. — Set  up  the  apparatus  shown  in  Fig. 
5  and  distill  some  water  in  it.  The  flask  should  be 
half  filled  with  water  which  has  been  colored  by  ink.  It 
must  be  perfectly  dry  on  the  outside,  and  is  heated  gently 
with  a  piece  of  wire  gauze  between  it  and  the  flame  to 
avoid  breakage.  Before  heating,  connect  the  condenser 
with  the  water  supply  and  pass  a  gentle  current  of  water 
in  at  the  lower  tube,  letting  it  run  off  at  the  upper  tube 
into  the  sink.  Avoid  violent  boiling,  which  would  carry 
over  the  impurities  into  the  receiver.  Distill  about  100 


WATER 


cubic  centimeters  1  of  water,  stopping  before  the  water  in 
the  flask  becomes  very  low,  otherwise  the  latter  would 
break.  Test  the  purity  of  the  water  by  evaporating  a 
small  quantity  in  a  perfectly  clean  dish. 


Fio.  5. 

How  does  distillation  remove  the  impurities  from 
water  ?  Can  all  impurities  be  removed  in  this  way  ?  Why 
is  glass  apparatus  not  used  in  making  distilled  water  com- 
mercially? Make  a  drawing  of  the  apparatus. 

EXPERIMENT  5. — Fill  a  beaker  with  crushed  ice  and 
clamp  a  thermometer  so  that  the  bulb  is  surrounded  by 
the  mass.  Notice  the  behavior  of  the  mercury  column 
and,  when  it  becomes  stationary,  take  several  readings. 
Take  the  average  of  the  temperatures  and  record  it  as 
the  result  of  a  determination  of  the  melting-point  of  ice. 

Suppose  you  had  placed  liquid  water  in  the  beaker; 
had  surrounded  the  beaker  with  a  freezing  mixture  and 
taken  the  temperature  when  the  water  was  partially  frozen, 
would  the  result  have  been  the  same  ? 

What  evidence  have  we  that  ice  and  water  are  two 
forms  of  the  same  substance  ? 

Make  a  mixture  of  salt  and  crushed  ice  and  take  the 
temperature  as  before.  (  ?) 

1  c.c.  will  be  used  hereafter  as  the  abbreviation  for  cubic  centi- 
meters.    100  c.c.  is  about  one-half  the  capacity  of  a  kitchen  cup. 


8 


ELEMENTARY  CHEMISTRY 


EXPERIMENT  6. — Half  fill  a  small  beaker  with  distilled 
water,  dry  it  on  the  outside  and  support  it  on  wire  gauze. 
Clamp  a  thermometer  so  that  the  bulb  is  in  the  liquid 
and  heat.  Watch  the  thermometer.  Small  bubbles  escape 
before  the  liquid  begins  to  boil.  Explain  this.  What 
would  be  the  effect  of  putting  fish  in  water  which  had  been 
boiled  and  then  cooled  in  the  absence  of  air? 

When  the  liquid  boils,  take  several  readings  of  the 
temperature.  Kecord  the  results.  What  is  the  effect  of 
variation  of  pressure  on  the  boiling-point  of  water.  Why 
is  it  difficult  to  cook  certain  kinds  of  food  on  high  moun- 
tain-tops ? 

With  a  pipette  measure  off  50  c.c.  of  water  into  a 
100  c.c.  beaker.  Dissolve  in  the  water  10  grams  of  salt 
and  take  the  boiling-point  of  the  solution.  Repeat,  using 
10  grams  of  sugar.  How  do  the  boiling-points  of  solu- 
tions of  solids  compare  with 
that  of  pure  water?  Does 
sugar  or  salt  more  strongly 
affect  the  boiling-point? 

EXPERIMENT  7.  —  Carry 
out  the  electrolysis  of  water 
in  the  apparatus  shown  in 
Fig.  6.  Dilute  some  sul- 
phuric acid  with  about  ten 
times  its  volume  of  water. 
In  diluting  sulphuric  acid 
pour  the  acid  into  the  water 
in  a  thin  stream,  stirring 
constantly — never  the  water 
into  the  acid.  (With  the  other  acids  it  is  a  matter  of  in- 
difference.) Cool  this  liquid  by  standing  it  in  a  tin  pan 
full  of  water.  Fill  the  apparatus  with  the  liquid  by  pour- 
ing it  into  the  funnel  tube  until  it  just  reaches  the  stop- 
cocks. The  latter  must  be  open  during  this  operation. 


FIG.  6. 


WATER  9 

Close  them  and  connect  the  two  pieces  of  platinum  foil 
in  the  apparatus  with  a  battery  or  a  dynamo-circuit.  Al- 
low the  current  to  pass  until  about  15  c.c.  of  the  gas 
which  is  formed  in  smaller  quantity  has  collected.  Do 
not  let  the  level  of  the  liquid  fall  as  far  as  the  plati- 
num foil. 

Stop  the  current,  and  if  the  apparatus  is  graduated, 
read  off  carefully  the  volumes  of  both  gases.    Read  from 
the  bottom  of  the  meniscus  (Fig.  7).    If  the  apparatus  is 
not  graduated,  measure  the  length  of 
each  gas-column  with  a  meter  scale. 
What  is  their  relation  by  volume? 

Cautiously  open  each  stop-cock  for 
an  instant  to  drive  out  any  water 
which  may  be  above  them.  Obtain  a 
pine  splint  bearing  a  spark  by 'extin- 
guishing the  flame  of  the  burning 
splint  by  a  quick  movement  of  the 
hand.  Let  the  gas  which  is  present  in 
smaller  quantity  stream  out  against 
the  spark.  (  ?)  Slip  over  the  end  of  the  other  tube  a  very 
short  piece  of  rubber  tubing  connected  with  a  short  piece 
of  glass  tubing  drawn  out  to  a  fine  jet.  Let  the  gas  escape 
against  the  flame  of  a  burning  match,  and  remove  the 
flame.  ( ?)  What  evidence  does  this  experiment  furnish 
of  the  composition  of  water?  Why  does  the  level  of  the 
liquid  in  the  apparatus  sink  so  slowly  when  it  is  run  with 
the  stop-cocks  open,  although  much  gas  escapes?  What 
becomes  of  the  sulphuric  acid  and  what  is  the  object  of 
adding  it?  (Part  I,  p.  6). 

EXPERIMENT  8. — Qualitative  synthesis  of  water. — Fill 
a  U-shaped  tube  with  fragments  of  dry  calcium  chloride, 
free  from  powder.  Cork  the  open  ends  of  the  tube  tightly 
and  connect  it  on  the  one  side  with  a  Kipp  generator  fur- 
nishing a  current  of  hydrogen,  and  on  the  other  with  a 


10 


ELEMENTARY  CHEMISTRY 


short  tube  of  hard  glass  drawn  out  to  a  jet  (Fig.  8).    Pass 
a  gentle  current  of  hydrogen  through  the  tube  for  a  few 


FIG.  8. 

seconds  to  drive  out  the  air,  and  then  light  the  gas  at  the 
jet.  Hold  over  the  flame  a  cold,  dry  funnel.  (?)  What  is 
the  conclusion?  Make  a  drawing 
of  the  apparatus.  The  object  of  the 
calcium  chloride  is  to  dry  the  gas. 
Why?  Hold  a  cold,  dry  cylinder 
or  bottle  over  a  burning  candle  for 
a  moment.  Over  a  small  gas  flame. 
(?)  What  evidence  does  this  fur- 
nish with  respect  to  the  composi- 
tion of  the  gas  and  the  candle  ? 

EXPERIMENT  9. — Quantitative 
synthesis  of  water.1 — This  experi- 
ment is  made  in  the  U-shaped 
eudiometer  (Fig.  9).  The  end  at 
0  is  open.  The  other  limb  of  the 


Fio.  9. 


1  This  experiment  requires  some  ex- 
pertness  in  manipulation.  It  will  be 
well  to  carry  it  out  on  the  lecture-table 
at  this  stage,  or  to  postpone  it  until  the 
student  has  acquired  some  skill  in  the 
handling  of  apparatus. 


WATER 


11 


U  is  closed  by  a  three-way  stop-cock  8.  This  stop-cock 
is  so  constructed  that  gas  passed  in  through  T  will  either 
escape  into  the  air  by  the  stem  E  or  pass  down  into  the 
eudiometer,  according  to  the  position  of  the  stop-cock. 
Take  out  the  stop-cock  and  inspect  it  until  you  understand 
why  this  should  be  the  case.  If  necessary  smear  a  little 
vaseline  thinly  over  it,  put  it  back  in  place  and  turn  it  un- 
til the  vaseline  is  evenly  distributed.  This  will  make  it  air- 
tight, which  is  essential  to  the  success  of  the  experiment. 

Open  8  and  close  the  stop-cock  at  C.  Pour  mercury 
in  at  0  until  the  apparatus  is  filled  to  the  level  of  E,  but 
not  above  it. 

Fit  a  test-tube  with  a  perforated  cork  and  delivery 
tube  (Fig.  10).  Over  the  end  of  the  delivery  tube  slip  a 
short  piece  of  rubber  tubing  which  will  readily  slip  over 
the  stem  T,  making  an  air-tight 
junction.  Fill  this  test-tube  to  the 
depth  of  2  cm.  (£  inch)  with  a 
mixture  of  1  part  manganese  diox- 
ide and  3  parts  potassium  chlorate. 
Heat  very  gently,  proceeding  exact- 
ly as  directed  in  Experiment  23. 
From  time  to  time  test  the  gas 
given  off  through  the  delivery  tube 
with  a  splint  bearing  a  spark. 
When  the  spark  bursts  into  flame, 
adjust  the  three-way  stop-cock  of 
the  eudiometer  so  that  gas  admitted 
at  T  will  escape  at  E,  and  let  oxygen  stream  through  the 
tip  and  stop-cock  for  ten  or  fifteen  seconds.  (Why  ?)  Then 
turn  the  stop-cock  so  that  the  oxygen  passes  into  the  tube 
over  the  mercury  and  allow  5  to  8  c.c.  of  oxygen  to  pass 
in.  Turn  the  stop-cock  so  that  the  oxygen  in  the  eudi- 
ometer is  shut  off  both  from  the  external  air  and  from 
the  oxygen-generator,  and  at  once  remove  the  latter. 


FIG.  10. 


12  ELEMENTARY   CHEMISTRY 

Place  a  dry  beaker  under  the  stop-cock  C,  and  allow  mer- 
cury to  run  out  until  the  level  of  the  mercury  in  the  two 
limbs  is  the  same.1  (Why?)  Be  careful  in  all  these 
manipulations  not  to  touch  with  the  hands  that  part  of 
the  tube  containing  the  oxygen.  (Why?)  Read  and 
record  the  volume  of  the  oxygen.  Repeat  the  reading. 

Now  transfer  your  cork  and  delivery  tube  to  another 
test-tube  in  which  you  have  placed  a  few  pieces  of  granu- 
lated zinc,  about  3  cm.  (1  inch)  of  water  over  them  and 
enough  sulphuric  acid — added  very  gradually — to  produce 
a  rather  brisk  evolution  of  hydrogen.2  Turn  the  stop-cock 
so  that  gas  passed  in  at  T  will  escape  at  E.  Be  careful 
in  doing  this  not  to  throw  the  oxygen  into  communication 
with  the  air,  as  this  would  spoil  the  experiment. 

Slip  the  rubber  tube  over  T  and  allow  the  hydrogen 
to  stream  out  through  E  for  about  two  minutes.  (Why?) 
Pass  about  three  times  as  much  hydrogen  as  you  have 
taken  of  oxygen  into  the  eudiometer,  turn  the  stop-cock 
so  as  to  shut  it  off  from  the  air  and  the  hydrogen  gener- 
ator, and  remove  the  latter.  Equalize  the  levels  of  the 
mercury  in  the  two  limbs,  read  off  the  volume  of  hydrogen 
added,  and  let  the  apparatus  stand  several  minutes  to  give 
the  two  gases  time  to  mix  thoroughly. 

The  two  gases  can  now  be  caused  to  combine  by  a  spark 
passed  between  the  platinum  wires  w  w.  The  spark  can 
be  obtained  from  an  induction  coil  excited  by  three  Edi- 
son-Lalande  or  by  two  dichromate  cells.  Since  the  ex- 
plosion may  possibly  be  violent  enough  to  wreck  the  ap- 
paratus, it  is  well  to  take  the  precaution  of  running  out 
mercury  at  C  until  the  level  of  the  mercury  in  that  limb 

1  If  by  mistake  too  much  mercury  is  run  out,  add  mercury  at  0 
until  the  level  is  higher  in  that  limb,  and  again  run  out  through  C 
until  the  levels  are  the  same. 

1  The  hydrogen  is  best  taken  from  a  Kipp  apparatus  if  one  is  at 
hand. 


WATER  13 

is  much  lower  than  in  the  other.  This  dilutes  the  gases 
and  reduces  the  violence  of  the  explosion.  Before  passing 
the  spark,  put  the  thumb  tightly  over  0  to  prevent  any 
mercury  being  thrown  out. 

Let  the  apparatus  stand  for  five  minutes,  equalize  the 
levels  of  the  two  mercury-columns  and  read  off  the  volume 
remaining.  Kecord  it.  Pour  mercury  into  0  and  show 
that  the  residual  gas  is  hydrogen  by  letting  it  escape  at  T 
and  burning  it.  Slip  a  short  rubber  tube  bearing  a  jet 
over  T  before  doing  this,  otherwise  the  heat  of  the  flame 
would  crack  the  apparatus.  Make  the  calculation  as  in 
the  following  example: 

Suppose  you  have  taken 7  c.c.  oxygen, 

and  that  the  volume  after  adding  hydrogen  is  27  c.c. ; 

then  the  volume  of  hydrogen  added  is ....  20  c.c. 

Suppose  further  that  the  volume  after  the  explosion  is 
6.1  c.c. 

This  residual  gas  is  shown  by  investigation  to  be  hy- 
drogen. It  is  clear  that  20  —  6.1  or  13.9  c.c.  of  hydrogen 
must  have  disappeared  with  the  7  c.c.  of  oxygen,  to  form 
water.  Hence  the  relation  by  volume  in  which  the  gases 
combine  is  7 :  13.9  or  1 :  1.98. 

Make  a  drawing  of  the  apparatus  in  your  note-book  and 
give  a  complete  description  of  your  work.  Explain  the  cal- 
culation. What  becomes  of  the  water  formed  in  the  explo- 
sion and  why  is  its  volume  not  considered  ?  Is  the  compo- 
sition of  water  always  the  same  ?  If  so,  how  do  you  account 
for  the  fact  that  different  experiments,  even  when  carried 
out  with  great  care,  always  yield  slightly  different  results  ? 

The  more  numerous  and  careful  the  experiments  the 
more  closely  the  average  result  of  them  will  approach  the 
ratio  1 :  2. 

What  general  statement  can  be  made  about  the  vol- 
umes in  all  cases  in  which  two  gases  combine? 


14  ELEMENTARY   CHEMISTRY 

How  can  you  calculate  the  composition  of  water  by 
weight  from  its  composition  by  volume? 

EXPERIMENT  10. — Action  of  oxygen  on  copper.  Syn- 
thesis of  water  from  hydrogen  and  copper  oxide. — Fit  up 
the  apparatus  shown  in  Fig.  11.  T  T'  is  a  piece  of  hard 


PIG.  11. 

glass  tubing,  1  cm.  or  more  in  diameter  and  about  30  cm. 
(1  foot)  long.  In  order  to  cut  off  a  suitable  piece  of 
tubing,  first  make  a  notch  on  the  tube  at  the  point  where 
it  is  desired  to  cut  it.  Then  wrap  several  layers  of  wet 
filter  paper  around  the  tube  quite  close  to  the  notch,  and  do 
this  also  on  the  other  side  of  the  notch  (Fig.  12),  so  that 

a  narrow  band  of 
glass,  1  cm.  wide  or 
less,  with  the  notch 
is  left  bare.  If  this 
bare  portion  is  held  in  the  burner  flame,  it  will  usually 
break  off  evenly,  and  any  projecting  portions  can  be  re- 
moved by  careful  chipping  with  a  key. 

Thrust  a  loose  plug  of  asbestos  into  the  tube  about 
5  cm.  (2  inches)  from  one  end,  and  upon  this  pour  a  col- 
umn of  copper  clippings  about  20  cm.  (8  inches)  long. 
Put  a  similar  loose  plug  of  asbestos  in  the  other  end  of 
the  tube  to  hold  the  clippings  in  place.  Holding  the  tube 
horizontal,  tap  it  gently  upon  the  table  so  as  to  cause  the 


FIG.  12. 


WATER  15 

clippings  to  settle,  and  make  a  channel  for  the  passage  of 
gases. 

Select  two  corks  a  trifle  too  large  to  fit  in  the  ends  of 
the  tube,  and  roll  them  carefully  under  the  foot  until  they 
fit.  Perforate  the  corks  either  with  a  rat-tail  file  or  a 
cork-borer,  and  fit  them  with  tubes,  as  in  the  cut.  In  per- 
forating with  a  rat-tail  file,  thrust  the  sharp  end  of  the 
file  through  the  middle  of  the  cork,  beginning  with  the  - 
small  end,  and  resting  the  cork  on  the  table.  Then  enlarge 
the  hole  with  the  rough  part  of  the  file  until  the  glass 
tube  fits  it  tightly.  The  glass  tubes  must  be  rounded  at 
the  ends,  and  must  be  introduced  with  a  twisting  motion 
and  without  strong  pressure,  the  use  of  which  is  likely 
to  break  the  tube  and  seriously  cut  the  hand. 

Clamp  the  tube,  or  suspend  it  by  a  wire,  and  apply  heat 
from  a  wing-top  burner,  very  gently  at  first.  If  clamps 
are  used,  they  must  be  placed  at  the  ends,  so  as  not  to  be 
burned.  Raise  the  temperature  gradually  until  the  copper 
is  at  a  dull  red-heat.  If  the  tube  softens  or  shows  signs 
of  bending,  moderate  the  heat. 

Now  connect  the  short  straight  glass  tube  with  a  test- 
tube  in  which  oxygen  is  being  generated  from  potassium 
chlorate  and  manganese  dioxide  (Experiment  23),  and 
pass  oxygen  through  it.  What  is  the  result?  Describe 
the  product.  What  is  its  composition?  At  which  end  of 
the  tube  does  the  change  begin,  and  why? 

Remove  the  oxygen  generator  and  connect  the  tube 
with  a  hydrogen  generator.  The  Kipp  apparatus  is  best. 
//  an  ordinary  generator  is  used,  hydrogen  must  be  allowed 
to  escape  briskly  from  the  exit  tube  for  at  least  five  min- 
utes before  connecting  it  to  the  red-hot  tube.  The  hydro- 
gen must  be  passed  through  a  U-shaped  tube  containing 
dry  calcium  chloride  in  lumps — not  powder — to  free  it 
from  water  before  it  enters  the  copper  tube.  See  Experi- 
ment 22  for  the  method  of  generating  hydrogen. 


16  ELEMENTARY   CHEMISTRY 

What  change  takes  place  in  the  black  substance?  Hy- 
drogen has  a  similar  effect  on  the  compounds  of  oxygen 
with  many  metals — for  instance,  upon  tin  oxide  and  lead 
oxide.  The  effect  is  called  reduction.  Some  other  oxides, 
like  those  of  magnesium  and  aluminium,  are  not  affected 
by  it. 

Look  in  the  cooler  part  of  the  tube,  and  in  the  bent 
tube,  for  evidence  of  the  formation  of  water.  What  evi- 
dence of  the  composition  of  water  is  furnished  by  this 
experiment  ? 

PKOBLEMS 

1.  The  electric  current  is  passed  through  water  containing 
sulphuric  acid  until  20   grams  of  the  water  have   disappeared. 
What  weights  of  hydrogen  and  of  oxygen  have  been  liberated  ? 

2.  300  grams  of   water  are  decomposed.     What  weights  of 
hydrogen  and  oxygen  are  produced  ? 

3.  If  27.396  grams  of  water  contain  24.352  grams  of  oxygen, 
what  is  the  percentage  composition  of  water  ? 

4.  A  quantity  of  hydrogen  gas  weighing  4  grams  is  slowly 
passed  through  a  glass  tube  containing  a  large  quantity  of  copper 
oxide  heated  to  redness,     (a)  How  much  will  this  tube  lose  in 
weight  ?     (ft)  What  weight  of  water  will  be  produced  ? 

5.  50  c.c.  of  hydrogen  are  mixed  with  50  c.c.  of  oxygen  and 
the  mixture  exploded,     (a)  If  the  process  takes  place  below  100°, 
which  gas  remains  in  the  tube,  and  how  much  ?     (b)  If  the  whole 
process  takes  place  above  100°,  what  is  the  total  volume  remain- 
ing in  the  tube,  and  of  what  does  it  consist  ? 


CHAPTEE   II 

SOLUTION 

EXPERIMENT  11.— Soluble  and  insoluble.— Place  a  lit- 
tle coarsely  powdered  rosin  in  a  test-tube,  cover  it  with 
water,  and  shake  it.  (?)  Pour  out  the  water,  cover  the 
rosin  with  alcohol,  and  shake.  ( ?)  Reserve  this  liquid. 


SOLUTION  17 

Shake  up  some  copper  sulphate,  first  with  alcohol,  and 
then  with  water.  (  ?)  Give  definitions  of  the  terms  solu- 
ble, insoluble,  and  solution. 

Pour  the  alcoholic  solution  of  rosin  into  a  large  quan- 
tity of  water  in  a  bottle  or  cylinder.  (?)  The  resulting 
state  of  things  is  called  a  suspension.  State  some  differ- 
ences between  solutions  and  suspensions. 

EXPERIMENT  12. — Saturated  solutions. — Half  fill  a 
liter  flask  with  water  and  add  small  quantities  of  salt, 
shaking  after  each  addition  until  the  liquid  will 
dissolve  no  more.  Cork  the  flask  and  shake  the 
liquid  containing  a  little  undissolved  salt  for  a 
long  time,  to  make  sure  that  it  will  not  take  up  any 
more.  Pour  off  the  liquid  into  a  narrow  cylinder 
and  take  its  density  by  means  of  a  hydrometer 
(Fig.  13).  Most  solutions  are  denser  than  water. 
How  about  salt-solution?  The  density  will  serve 
as  an  indication  of  the  amount  dissolved.  Of 
course  if  two  solutions  of  table-salt  have  the  same 
density,  they  contain  the  same  percentage  of  dis- 
solved salt. 

Half  fill  the  flask  again  with  water  and  add  a 
large  quantity  of  salt,  several  times  as  much  as  the  water 
can  dissolve.  Shake  for  some  time.  The  liquid  should 
contain  a  sediment  of  undissolved  salt  more  than  5  cm.  in 
depth.  Finally,  let  the  liquid  settle,  pour  off  into  a  narrow 
cylinder,  and  take  the  density  as  before.  The  result  should 
be  the  same. 

This  important  experiment  is  an  example  of  the  fact 
that  the  quantity  of  a  substance  which  dissolves  to  produce 
a  saturated  solution  does  not  depend  upon  the  excess  of 
undissolved  solid  present  in  the  liquid.  There  must  be 
some  solid  present;  otherwise  the  liquid  would  not  be 
saturated;  but  a  solution  which  remains  in  contact  with  a 
small  quantity  of  a  solid  without  dissolving  any  more  of 


18  ELEMENTARY  CHEMISTRY 

it,  is  just  as  strong  as  one  which  is  in  contact  with  a 
large  quantity. 

EXPERIMENT  13. — Effect  of  temperature  on  solubility. 
Crystallization. — Powder  some  potassium  chlorate  in  a 
mortar,  not  by  pounding  it,  but  by  moving  the  pestle  with 
a  strong,  steady  pressure.  Place  50  c.c.  of  water  in  a 
100  c.c.  beaker  and  add  potassium  chlorate  to  it  in  small 
quantities.  Stir  until  each  portion  is  dissolved  before  add- 
ing the  next. 

When  the  liquid  is  saturated,  place  the  beaker  on  wire- 
gauze  on  a  ring  of  your  stand  and  heat  it  gently.  Continue 
stirring  and  adding  more  potassium  chlorate.  What  is 
the  effect  of  temperature  on  solubility?  Stand  the  beaker 
aside  to  cool.  (?)  What  is  the  most  obvious  difference 
between  a  crystal  and  a  bit  of  non-crystalline  matter  ? 

This  is  by  no  means  the  only  difference.  The  great  distinc- 
tion is  that  all  properties  which  have  direction  at  all,  are  different  in 
different  directions  in  a  crystal.  Thus,  a  crystal  breaks  more  eas- 
ily in  some  directions  than  in  others  ;  it  conducts  heat  better  in 
certain  directions  ;  it  transmits  light  quite  differently  along  cer- 
tain lines,  and  so  on.  Even  the  solubility  of  a  solid  is  different  on 
different  faces  of  a  crystal. 

For  all  these  reasons,  if  a  crystal  is  broken  into  any  irregular 
shape,  it  is  easy,  in  spite  of  this,  to  identify  it  as  a  piece  of  crys- 
talline matter ;  while  on  the  other  hand,  if  a  mass  of  glass  is  cut 
into  the  exact  shape  of  a  crystal,  it  is  equally  easy  to  show  that  it 
is  a  fragment  of  non-crystalline  (amorphous)  material  artificially 
shaped. 

Heat  some  clear  lime-water  to  boiling  in  a  test-tube. 
Take  the  lime-water  out  of  the  bottle  with  a  glass  tube, 
and  avoid  stirring  up  the  material  at  the  bottom.  This 
applies  only  to  lime-water.  All  ordinary  liquids  are  to 
be  poured  carefully  from  the  bottle,  not  withdrawn  by 
means  of  tubes.  Cork  the  lime-water  bottle  at  once  after 
using  it,  since  the  air  spoils  it.  Does  the  liquid  become 


SOLUTION  19 

turbid?  Does  heat  affect  the  solubility  of  lime  in  the 
same  way  as  it  does  that  of  potassium  chlorate?  How  do 
almost  all  solids  behave  in  this  respect? 

EXPERIMENT  14. — Supersaturated  solutions. — Place  a 
few  drops  of  water  in  a  test-tube,  half  fill  the  tube  with 
sodium  thiosulphate — called  "  hypo  "  by  the  photographer 
— and  heat  with  a  small  flame  kept  in  constant  motion. 
When  complete  solution  has  occurred,  pour  the  liquid  into 
a  clean  tube  and  cork  it  to  exclude  dust.  Let  it 
cool.  Does  it  behave  like  potassium  chlorate  solu- 
tion on  cooling?  Throw  into  the  cold  liquid  a 
fragment  of  solid  sodium  thiosulphate.  (?) 
What  is  a  supersaturated  solution?  Do  you  re- 
gard it  as  a  stable  or  an  unstable  state  of  things? 
Would  any  crystal  answer  the  purpose,  if  dropped 
into  the  supersaturated  solution  in  this  experi- 
ment? Since  a  solution  which  does  not  contain 
any  of  the  undissolved  solid  may  be  either  un- 
saturated  or  supersaturated,  what  is  the  only  way 
of  being  sure  that  a  solution  is  saturated? 

EXPERIMENT     15. — Solutions    of    liquids    in   ] 
liquids. — To  a  test-tube  of  water  add  a  drop  of  chloro- 
form.    Cover  the  tube  with  the  thumb  and  shake  it.     Are 
the  liquids  apparently  insoluble  in  each  other?    Are  they 
really  insoluble  in  each  other  ? 

Half  fill  a  small  separatory  funnel  (Fig.  14)  with  water. 
Add  a  little  ether,  insert  the  stopper  and  shake.  Ether  is 
highly  inflammable,  and  must  not  be  used  in  the  neighbor- 
hood of  a  flame.  The  ether-bottle  must  be  kept  tightly 
corked.  Add  another  small  quantity  of  ether  and  shake 
again.  (?)  Now  add  about  half  as  much  ether  as  there  is 
water  present,  and  shake.  Is  ether  soluble  in  water?  Is 
there  a  limit  to  its  solubility?  What  is  the  composition 
of  the  two  layers  in  the  funnel?  Give  a  suitable  name 
to  each  layer.  Allow  about  30  c.c.  of  the  lower  layer  to 
24 


20  ELEMENTARY  CHEMISTRY 

run  into  a  100  c.c.  beaker,  and  get  rid  of  the  remainder  of 
the  liquid  by  running  it  down  the  sink  with  an  abundance 
of  water.  Place  the  beaker  on  wire  gauze  and,  having  first 
made  sure  that  no  one  in  the  neighborhood  is  working  with 
ether,  heat  it  gently.  Hold  a  lighted  match  over  the 
beaker.  (?)  Did  the  lower  layer  contain  ether?  Is  the 
solubility  of  ether  in  water  increased  or  decreased  by  heat- 
ing ?  The  solubility  of  water  in  ether  is  increased  by  heat. 

Fill  a  test-tube  one-fourth  with  alcohol  and  add  water 
in  small  quantities,  shaking  after  each  addition  until  the 
tube  is  full.  What  is  the  result?  Is  it  possible  to  obtain 
two  layers  by  mixing  alcohol  and  water?  Some  other 
pairs  of  liquids — for  instance,  alcohol  and  ether — behave 
in  the  same  way. 

EXPERIMENT  16. — Solutions  of  gases  in  liquids. — Heat 
some  faucet  water  in  a  beaker  on  wire  gauze  not  quite  to 
boiling.  (?)  Heat  some  ammonia  water  in  a  beaker.  Hold 
a  burner  flame  over  the  beaker.  Is  there  any  evidence  that 
a  gas  escapes?  What  is  the  effect  of  rising  temperature 
on  the  solubility  of  a  gas? 

Fill  a  100  c.c.  beaker  with  soda-water  from  a  siphon 
bottle.  What  is  soda-water?  What  was  the  cause  of  the 
effervescence  when  the  liquid  escaped  into  the  beaker? 
How  does  the  solubility  of  a  gas  vary  with  the  pressure? 
Lower  into  the  liquid  a  sealed  tube,  open  end  down. 
Throw  into  it  a  fragment  of  charcoal.  Explain  the 
results. 

Taste  some  of  the  soda-water.  Place  about  10  c.c.  of 
it  in  a  test-tube  and  boil  it  for  a  time.  Cool  by  running 
water  over  the  tube  and  taste  again.  Is  the  taste  the  same  ? 
Why  ?  What  is  the  cause  of  the  unpleasant  taste  of  boiled 
water,  and  how  can  it  be  remedied  ? 


PHYSICAL  AND  CHEMICAL  CHANGE  21 


PROBLEMS 

6.  10.98  grams  of  a  solution  of  potassium  chlorate  saturated  at 
18°  was  placed  in  a  weighed  dish  and  evaporated  to  dryness.    The 
residue  weighed  .7025  gram.     How  much  potassium  chlorate  was 
contained  in  100  parts  of  the  solution  ? 

7.  Taking  the  figures  stated  in  problem  6,  how  much  potas- 
sium chlorate  will  100  parts  of  water  at  18°  dissolve  ? 

8.  Taking  the  same  figures,  how  much  water  at  18°  is  neces- 
sary to  dissolve  1  gram  of  potassium  chlorate  ? 

9.  A  solution  of  common  salt  saturated  at  15°  contained  26.39 
per  cent  of  salt.     How  much  salt  will  100  grams  of  water  dissolve 
at  15°  ? 


CHAPTER    III 

PHYSICAL  AND    CHEMICAL    CHANGE-LAW  OF  THE 
INDESTRUCTIBILITY   OF  MATTER 

EXPERIMENT  17. — Changes  in  matter. — Take  a  piece 
of  platinum  foil  in  forceps  and  hold  in  the  Bunsen  flame, 
Let  it  cool.  (  ?) 

Tear  up  a  little  paper  in  very  small  pieces  and  bring 
near  it  a  roll  of  sulphur.  Rub  the  sulphur  briskly  on  the 
coat-sleeve  and  again  bring  it  near  the  paper.  (?) 

Crease  a  long  narrow  piece  of  paper  in  the  middle  so 
as  to  make  a  trough,  which  will  slip  into  a  test-tube.  By 
means  of  this  introduce  a  little  mercuric  iodide  into  a  clean, 
dry  test-tube,  without  getting  any  of  it  on  the  sides.  In 
doing  this  hold  the  tube  horizontally  and  slip  it  over  the 
paper,  near  the  end  of  which  the  mercuric  iodide  is  placed. 
Then  upset  the  trough  and  deposit  the  mercuric  iodide  in 
the  bottom  of  the  tube.  Always  use  this  method  of  intro- 
ducing powders  into  tubes  when  you  desire  to  keep  the 
upper  portion  of  the  latter  clean. 

Lay  a  cork  loosely  in  the  mouth  of  the  tube  and  weigh 


22  ELEMENTARY  CHEMISTRY 

it  or  balance  it  with  copper  filings  or  iron  filings.  Before 
weighing  be  sure  that  it  is  absolutely  clean  and  dry  on 
the  outside,  and  all  through  the  experiment  handle  the 
tube  with  paper — not  with  the  fingers— to  avoid  soiling  it. 

Now  heat  the  tube  very  gently.  It  should  be  simply 
brushed  with  the  flame  once  in  two  or  three  seconds  and 
the  powder  should  be  constantly  shaken  about.  The  tem- 
perature must  not  rise  much  above  the  boiling-point  of 
water.  Overheating  will  melt  the  mercuric  iodide  and 
spoil  the  result. 

When  the  change  is  complete  stop  heating,  let  the  tube 
cool,  and  weigh  it  again.  Is  the  weight  the  same?  What 
does  the  result  show  ?  What  is  the  substance  in  the  tube  ? 
Throw  out  half  of  it  on  a  piece  of  paper  and  rub  it  with 
a  glass  rod.  (?)  Let  the  rest  of  it  stand  over  night  and 
examine  it  in  the  morning.  (?)  What  is  the  natural  state 
of  mercuric  iodide  at  ordinary  temperatures  ?  At  slightly 
elevated  temperatures  ?  In  what  respects  does  this  change 
resemble  and  in  what  respects  differ  from  the  transforma- 
tion of  water  into  ice  ?  In  what  respect  do  all  the  changes 
carried  out  in  Experiment  17  resemble  each  other?  Men- 
tion some  other  changes  in  matter  which  are  like  them  in 
this  respect. 

EXPERIMENT  18. — Changes  in  matter.  The  law  of  the 
indestructibility  of  matter. — Take  a  short  piece  of  mag- 
nesium ribbon  in  forceps  and  hold  one  end  of  it  in  the 
flame.  Receive  the  product  in  a  dish  and  examine  it. 
It  is  called  magnesium  oxide.  Why?  Compare  this 
change  with  the  heating  of  platinum. 

Place  in  a  small  hard  glass  test-tube  enough  mercuric 
oxide  to  fill  it  to  the  depth  of  1  cm.  or  more.  Introduce 
the  powder  by  means  of  a  paper  trough.  Clamp  the  tube 
horizontally,  placing  the  clamp  near  the  mouth  so  that 
heat  can  be  applied  without  spoiling  the  clamp,  and  heat 
gently,  brushing  the  tube  with  a  small  flame.  Hard-glass 


PHYSICAL  AND   CHEMICAL  CHANGE 


23 


tubes  must  be  heated  with  great  caution,  for  they  break 
very  readily.  When  the  mercuric  oxide  changes  color, 
examine  it  and  let  the  tube  cool..  What  kind  of  a  change 
is  this  ?  When  mercuric  acid'  is  cooled  to  — 200°  C.  it 
becomes  sulphur-yellow,  and  resumes  its  original  tint  on 
being  allowed  to  warm  to  room  temperature. 

Fit  the  test-tube  with  a  perforated  cork  bearing  a 
delivery  tube,  and  arrange  the  apparatus  as  shown  in  Fig. 
15.  Apply  heat,  at  first  very  gently,  gradually  raising  the 


FIG.  15. 

temperature  to  redness.  Collect  two  test-tubes  and  a  small 
cylinder  of  the  gas  given  off.  Apply  the  yellow  flame  of 
the  burner  and  cover  the  hot  tube  with  a  layer  of  soot  to 
make  it  cool  gradually  and  avoid  cracking.  Disconnect 
the  apparatus.  Examine  the  residue  in  the  tube.  What  is 
it?  Record  the  properties  of  the  gas  collected.  Into  a 
test-tube  of  it  introduce  a  spark.  ( ?)  Slip  a  glass  plate 
under  the  cylinder  full  of  the  gas  and  place  it  upright  on 
the  desk.  Draw  aside  the  plate  for  an  instant  and  intro- 
duce into  the  gas  a  piece  of  burning  candle,  held  erect  on 
a  bent  piece  of  stiff  wire.  (?)  What  is  the  gas?  Keep 
the  flame  out  of  contact  with  the  glass. 

Sketch  the  apparatus  in  your  note-book  and  write  a 
full  description  of  your  work.  What  is  the  historical  in- 
terest of  this  experiment  ?  Clean  the  hard-glass  tube  when 


24  ELEMENTARY  CHEMISTRY 

perfectly  cold  by  washing  it  out  thoroughly  with  water, 
then  allowing  a  little  strong  nitric  acid  to  remain  in  it 
for  a  time  and  finally  washing  with  water  again. 

Weigh  roughly  about  1  gram  of  potassium  iodide  and 
dissolve  it  in  a  small  test-tube  about  half  full  of  water. 
Weigh  roughly  0.7  gram  of  mercuric  chloride  and  dissolve 
it  in  about  100  c.c.  of  water  in  an  Erlen- 
myer  flask.  Mercuric  chloride  is  very 
corrosive  and  poisonous,  and  must  not  be 
touched  with  the  fingers.  If  any  of  it  gets 
upon  the  hands  it  must  be  removed  at  once 
by  washing  with  water.  Carefully  place 
the  test-tube  in  the  flask.  Remember 
that  chemical  glassware  is  purposely  made 
ig  thin  and  will  not  endure  violent  treatment. 
Hold  the  flask  upright,  see  that  it  is  ab- 
solutely dry  and  clean  on  the  outside,  and  cork  it  tightly. 
The  completed  apparatus  is  shown  in  Fig.  16. 

Now  take  hold  of  the  flask  by  means  of  a  folded  paper 
and  place  it  on  the  trip-scales  or  a  large  lecture-table  bal- 
ance. Weigh  it  accurately,  or  balance  it  by  metal  filings. 
Record  the  weight,  or  leave  the  filings  on  the  balance ;  take 
the  flask  with  the  paper  holder  and  tilt  it  so  that  the 
liquids  mix.  What  happens?  The  red  substance  is  mer- 
curic iodide.  This  accounts  for  the  mercury  and  the  iodine, 
but  there  must  be  another  compound  produced.  Of  what 
two  elements  ?  What  becomes  of  this  compound  ?  Devise 
a  method  of  obtaining  it. 

Put  the  flask  back  on  the  balance  and  weigh  it  again. 
How  do  the  weights  compare?  What  law  have  you  exem- 
plified? State  the  law  in  several  different  ways.  Why  is 
it  important?  What  evidence  does  astronomy  furnish  of 
its  truth? 


MIXTURE-ELEMENT-COMPOUND  25 

CHAPTER    IV 

MIXTURE-ELEMENT-COMPOUND— THE   LAW   OF  DEFINITE 
PROPORTIONS 

EXPERIMENT  19. — Mixtures  and  compounds. — Powder 
finely  some  roll  sulphur  (about  3  grams)  and  mix  most  of 
it  with  about  an  equal  weight  of  iron  filings.  Reserve  a 
little.  Examine  a  little  of  the  mixture  with  a  lens.  Can 
you  perceive  iron  and  sulphur  in  it?  Place  a  little  of  the 
mixture  in  a  beaker  half  full  of  water  and  stir  it.  (?) 
Devise  a  method  of  separating  the  two  based  upon  this 
behavior. 

Place  a  little  iron  filings  in  one  dry  test-tube,  and  a 
little  powdered  roll  sulphur  in  another,  and  try  the  beha- 
vior of  both  with  a  few  drops  of  carbon  disulphide.  Which 
dissolves  ? 

Carbon  disulphide  is  inflammable  and  must  not  be  used 
in  the  vicinity  of  flame.  .The  bottle  must  be  kept  corked. 
The  vapor  is  injurious  and  should  be  inhaled  as  little  as 
possible. 

Now  place  a  little  of  the  mixture  of  iron  and  sulphur 
in  a  dry  test-tube,  and  shake  it  up  with  about  5  c.c.  of  car- 
bon disulphide.  Allow  to  settle,  and  pour  off  the  clear 
liquid  into  a  dry  dish.  What  is  the  substance  in  the  test- 
tube?  What  remains  in  the  dish  after  the  carbon  disul- 
phide has  evaporated? 

Place  a  small  portion  of  the  mixture  of  iron  and  sul- 
phur on  a  piece  of  paper  and  stir  it  with  a  magnet.  Ex- 
plain. 

Now  transfer  the  rest  of  the  mixture  to  a  dry  test-tube, 
tap  the  tube  on  the  table  to  make  a  channel  along  the 
upper  surface,  clamp  the  tube  horizontally  near  the  mouth 
and  apply  a  burner  flame  to  the  extreme  bottom.  It  is  well 
to  place  a  tin  pan  under  the  tube  in  case  of  breakage. 


26  ELEMENTARY   CHEMISTRY 


j  and  in  all  such  cases,  do  not  keep  the  hand  directly 
under  the  tube  while  healing.  When  the  reaction  begins, 
remove  the  flame  and  observe.  Is  energy  evolved  or  ab- 
sorbed when  iron  and  sulphur  combine  ?  Let  the  tube  cool, 
break  it,  if  necessary,  and  examine  the  product.  It  is  a 
chemical  compound  of  iron  and  sulphur,  called  iron  sul- 
phide. Powder  it,  and  examine  some  of  it  with  a  lens. 
Are  iron  and  sulphur  visible  in  it?  Will  carbon  disul- 
phide  extract  the  sulphur  from  it? 

Explain  all  the  distinctions  between  a  compound  and 
a  mixture  which  are  exemplified  in  this  experiment.  Notice 
that  all  these  differences  are  the  result  of  the  fact  that  the 
mixture  is  composed  of  two  substances,  each  having  its 
own  properties,  while  the  compound  is  one  substance. 

However,  the  greatest  distinction  between  compounds 
and  mixtures  does  not  appear  at  all  in  this  experiment. 
What  is  it? 

The  exact  quantities  in  which  iron  and  sulphur  com- 
bine to  form  iron  sulphide  is  56  parts  of  iron  to  32  parts 
of  sulphur.  Clearly  we  took  more  sulphur  than  was  neces- 
sary. What  became  of  it?  What  would  have  been  left 
in  the  test-tube  if  we  had  taken  more  iron  than  was  re- 
quired for  the  sulphur? 

EXPERIMENT  20. — The  law  of  definite  proportions. — 
Clean  and  dry  a  porcelain  crucible  carefully  and  weigh  it 
with  the  cover.  Record  the  weight.  Count  the  weights 
at  least  twice,  once  while  they  are  upon  the  pan  of  the 
balance  and  once  upon  returning  them  to  the  box,  to  be 
sure  there  is  no  mistake.  Handle  the  weights  with  for- 
ceps. They  must  not  be  touched  with  the  fingers.  Be 
sure  that  both  pans  of  the  balance  are  clean  before 
weighing. 

Introduce  into  the  crucible  from  0.3  to  0.5  of  a  gram 
of  powdered  magnesium  and  weigh  again  accurately.  Be 
sure  that  none  of  the  magnesium  is  upon  the  balance  pans, 


MIXTURE— ELEMENT— COMPOUND  27 

since  this  will  make  all  your  work  useless.     Again  count 
the  weights  at  least  twice. 

Support  the  covered  crucible  on  a  pipe-stem  triangle 
on  a  ring  of  the  stand  and  allow  a  Bunsen  flame  to  touch 
the  bottom.  After  ten  minutes  place  the  cover  so  as  to 
allow  a  very  limited  access  of  air.  The  operation  must  be 
so  conducted  that  nothing  escapes  from  the  crucible.  If 
any  white  smoke  (magnesium  oxide)  should  appear,  the 
crucible  must  be  covered  at  once.  After  ten  minutes  more, 
lift  the  cover  by  the  ring  with  the  forceps  and  shake  into 
the  crucible  any  magnesium  oxide  which  may  adhere  to  it, 
turn  up  the  flame  until  it  half  covers  the  open  crucible, 
and  heat  strongly  for  five  minutes,  turn  down  the  flame 
gradually,  extinguish  it,  cover  the  crucible,  let  cool  five 
minutes,  and  weigh.  The  magnesium  oxide  in  the  crucible 
should  be  white  or  grayish-white. 

The  following  will  serve  as  an  example  of  the  proper 
method  of  record  and  calculation: 
Weight  crucible,  cover,  and  magnesium 15.06 

"  "        and  cover,  empty 14.64 

Magnesium  taken 42 

Weight  crucible,  cover,  and  magnesium  oxide 15.34 

"  "  "      and  magnesium 15.06 

Oxygen  taken  from  air 28 

Therefore,  0.42  gram  of  magnesium  combines  with  0.28 
gram  of  oxygen.  Calculate  the  oxygen  which  will  com- 
bine with  1  gram  of  magnesium,  thus: 

.42  :  .28  =  1  :  x. 
x  =  $  gram,  or  .66. 

Calculate  also  the  quantity  of  magnesium  which  would 
combine  with  1  gram  of  oxygen,  thus: 
.28  :  .42  =  1  :  x. 
x  =  1.5  grams. 


28  ELEMENTARY  CHEMISTRY 

Calculate  in  the  same  way  the  quantity  of  oxygen  which 
would  combine  with  24  grams  of  magnesium  and  the  quan- 
tity of  magnesium  which  would  combine  with  16  grams 
of  oxygen. 

If  time  permits,  repeat  the  entire  experiment.  If  not, 
compare  your  results  with  those  of  others  who  have  done 
the  same  work.  If  the  experiment  is  properly  performed, 
different  results  should  be  nearly  the  same.  (Why  not 
exactly  the  same?)  If  your  result  varies  widely  from 
the  figures  given  above,  first  look  for  errors  in  the  calcula- 
tion, and  if  none  are  found,  repeat  the  experiment. 

Does  the  same  weight  of  magnesium  always  combine 
with  the  same  weight  of  oxygen?  Is  the  composition  of 
magnesium  oxide  always  the  same  ?  How  about  other  com- 
pounds? State  the  law.  What  bearing  have  the  results 
obtained  in  the  electrolysis  and  synthesis  of  water  upon 
the  law  ?  Explain  exactly  what  is  meant  by  the  statement 
that  magnesium  and  oxygen  are  elements  and  magnesium 
oxide  a  compound  of  them.  Why  are  solutions  mixtures 
although  they  are  homogeneous  and  can  not  be  separated 
by  mechanical  methods  ? 


CHAPTER   V 

HYDROGEN 

EXPERIMENT  21. — Reaction  of  sodium  and  water. — 

Sodium  must  not  be  touched  with  fingers  which  are  in  the 
least  moist  either  with  water  or  perspiration.  Your  desk 
must  be  dry  when  working  with  it  and  everything  with 
which  you  touch  it  must  be  scrupulously  dry.  The  air  acts 
rapidly  upon  sodium  and  it  must  not  be  exposed.  Take 
only  a  small  piece  from  the  bottle  at  a  time,  and  imme- 


HYDROGEN  29 

diately  cork  the  latter.  The  liquid  over  the  sodium  in  the 
bottle  is  naphtha  or  kerosene,  and  the  bottle  must  not  be 
opened  in  the  vicinity  of  a  flame.  No  sodium  must  be  put 
away  under  the  desk  or  allowed  to  remain  lying  about, 
since  it  may  catch  fire.  None  must  be  thrown  into  the 
waste- jar,  since  it  may  ignite  the  paper  or  other  substances 
which  the  jar  contains.  Return  any  unused  portions  to 
the  bottle,  or  place  in  a  vessel  specially  provided  for  that 
purpose. 

Throw  a  clean  piece  of  sodium,  free  from  crust,  half 
the  size  of  a  pea — no -larger — into  a  cylinder  half  full  of 
water.  Immediately  cover  the  cylinder  with  a  glass  plate 
or  a  piece  of  paper,  and  wait  until  the  reaction  is  over  be- 
fore removing  the  cover.  The  action  usually  ends  with  a 
slight  explosion  which  may  endanger  the  eyes  if  the  cover 
is  removed  too  soon. 

Describe  what  happens.  Throw  in  another  similar 
piece  of  sodium.  Feel  the  water  of  the  cylinder  between 
the  fingers.  ( ?)  Taste  a  little  of  it.  ( ?)  Here,  and  al- 
ways, immediately  reject  the  liquid  tasted  and  rinse  out 
the  mouth  with  water.  Do  not  taste  substances  unless 
directed.  Try  the  behavior  of  the  liquid  with  red  and  blue 
litmus  paper,  pieces  about  1  cm.  square  or  less.  (?)  Allow 
to  fall  into  the  liquid  a  drop  of  a  solution  of  phenol  phtha- 
lein.  (?)  What  are  these  new  properties  of  the  liquid 
due  to  ? 

Wrap  a  clean  piece  of  sodium  half  the  size  of  a  pea  in 
dry  tea-lead  (lead  foil).  Punch  several  holes  in  the  lead 
with  a  knife-blade.  Invert  a  test-tube  full  of  water  in 
a  tin  pan  containing  water,  and  quickly  slip  the  lead  con- 
taining the  sodium  under  it.  If  necessary  use  another 
smaller  piece  of  sodium  wrapped  in  lead  to  complete  the 
filling  of  the  tube.  Take  no  more  sodium  than  is  directed. 
The  use  of  larger  quantities  is  likely  to  cause  explosions 
which  may  imperil  the  sight. 


30  ELEMENTARY  CHEMISTRY 

What  gas  collects  in  the  tube  ?  Does  it  come  from  the 
water  or  the  sodium?  Has  the  gas  any  color  or  odor? 
Is  it  soluble  in  water?  Does  it  burn?  Quietly  or  with 
explosion?  To  what  product? 

Eecord  exactly  what  you  observe.  For  instance,  if  you 
find  the  gas  has  an  odor,  record  it,  and  if  you  think  the 
fact  peculiar,  inquire  of  the  instructor.  Never  alter  your 
observations  to  correspond  to  preconceived  notions. 

EXPERIMENT  22. — Hydrogen  from  zinc  and  sulphuric 
acid. — Place  about  20  grams  of  granulated  zinc  in  a  gas- 
generating  bottle.  The  bottle  must  be  held  almost  hori- 
zontal and  the  zinc  allowed  to  slide  into  it,  otherwise .  the 
shock  of  the  falling  zinc  will  break  it.  The  bottle  is  pro- 
vided with  a  doubly  perforated  cork  carrying  a  funnel  tube 
and  a  delivery  tube,  and  the  apparatus  is  arranged  for  col- 
lecting the  gas  over  water  (Fig.  10,  Part  I). 

Insert  the  cork  tightly  with  a  twisting  motion,  and 
pour  in  through  the  funnel  tube  enough  water  to  cover 
the  zinc  thoroughly.  The  funnel  tube  must  dip  into  this 
water.  Extinguish  any  burner  flame  that  may  be  in  the 
neighborhood,  and  slowly  add  strong  sulphuric  acid 
through  the  funnel  tube  until  gas  is  briskly  evolved.  Do 
not  add  too  much  acid.  The  maximum  should  be  about 
one-fourth  as  much  by  volume  as  there  is  water  present. 
Allow  the  gas  to  escape  through  the  water  for  three  min- 
utes. Why?  Do  not  attempt  at  any  time  in  the  experi- 
ment to  light  the  gas  at  the  exit  tube. 

Collect  the  gas  over  water  in  wide-mouthed  bottles  of 
about  300  c.c.  capacity.  Determine  its  properties.  Has 
it  any  color  or  odor?  Is  it  soluble  in  water?  Does  the 
method  of  collecting  it  throw  any  light  on  this  last  ques- 
tion? In  order  to  obtain  more  definite  information  fill 
a  test-tube  half  full  of  hydrogen  and  mark  the  level  of 
the  water  by  a  s  i  of  gummed  label  on  the  tube.  Then 
shake  the  tube  ^i  a  time,  keeping  its  mouth  under 


HYDROGEN  31 

water.  (?)  Use  this  method  hereafter  in  testing  the 
solubility  of  gases. 

Will  the  gas  burn?  Try  a  bottle  of  it — not  the  exit 
lube  of  the  generator.  Will  it  support  combustion?  Try 
it  by  holding  a  bottle-full  mouth  downward  and  intro- 
ducing a  lighted  candle  fastened  on  a  wire.  Keep  the 
candle  out  of  contact  with  the  walls  of  the  cylinder  so 
as  not  to  wet  the  wick.  Withdraw  the  candle  slowly.  (?) 
Repeat. 

Fill  two  bottles  of  the  same  size  with  the  gas.  Support 
one  in  an  inverted  position  in  a  ring  of  your  stand,  the 
mouth  not  touching  the  table.  Place  the  other  upright. 
Uncover  them  at  the  same  moment  and  allow  both  to  re- 
main uncovered  for  two  minutes.  Now  thrust  a  lighted 
taper  or  match  into  each  in  turn.  Draw  conclusions. 

Fill  a  small  strong  bottle  over  water  £  with  air  and  f 
with  hydrogen.  Ignite  the  mixture.  Explain  the  cause  of 
this  behavior?  Why  does  not  the  flame  strike  back  from 
gas-jets  along  the  mains  to  the  gas-works?  Would  it  be 
safe  to  supply  cities  with  a  mixture  of  gas  and  air  by 
means  of  pipes? 

Procure  a  strong  round-bottomed  ginger-ale  bottle  and 
a  cork  which  will  fit  it.  Perforate  the  cork  so  that  an 
artificial  straw  like  those  used  at  soda-fountains  will  fit 
fairly  well  in  the  hole.  Put  in  the  bottle  some  granulated 
zinc  and  water  and  add  enough  sulphuric  acid  to  produce 
effervescence.  Immediately  insert  the  cork  bearing  the 
straw  and  light  the  end  of  the  straw,  holding  the  bottle 
upright,  and  the  face  away  from  over  it.  What  is  the 
cause  of  the  explosion?  Would  an  explosion  have  oc- 
curred if  you  had  waited  several  minutes  before  lighting 
the  straw?  Why  is  it  dangerous  to  light  the  hydrogen 
at  the  exit  tube  of  your  generator  ?  Whenever  it  is  neces- 
sary to  do  this  you  should  wrap  the  generator,  cork,  and 
funnel-tube  carefully  with  a  towel,  so  that  no  harm  can 


ELEMENTARY   CHEMISTRY 


result,  and  wait  five  minutes  before  applying  the  light. 
This  applies  not  only  to  hydrogen  but  to  all  combustible 
gases. 

Filter  the  liquid  which  remains  in  your  gas-generating 
bottle  into  a  clean  porcelain  dish.     Prepare  the  filter  by 


FIG.  17. 

folding  it  first  into  a  semicircle,  then  into  a  quadrant 
(Fig.  17).  Thrust  it  tightly  into  a  funnel  and  wet  it 
to  keep  it  in  place.  Support  the  funnel  with  its  stem 

against  the  side  of  the 
dish,  so  that  the  liquid 
will  slip  down  the  side 
of  the  dish  and  not 
splash.  Pour  the  liq- 
uid into  the  filter  by 
means  of  a  glass  rod 
to  avoid  splashing 
(Fig.  18).  Never  fill 
up  quite  to  the  edge 
of  the  filter,  but  keep 
the  latter  nearly  full, 
as  this  makes  the 
process  more  rapid. 

When  the  liquid 
has  all  run  through 
wash  off  the  zinc  in  the  bottle,  and  return  it  to  the  stock 
bottle  of  granulated  zinc.  Evaporate  the  liquid  in  the  dish 
down  to  one-fourth  its  bulk.  No  wire  gauze  is  needed  in 


Pie.  18. 


OXYGEN  AND  HYDROGEN  PEROXIDE  33 

heating  the  dish.  The  -flame  must  not  be  allowed  to  touch 
that  portion  of  the  dish  which  is  above  the  level  of  the 
liquid  in  it,  and  the  dish  must  be  dry  on  the  outside.  The 
last  sentence  applies  to  all  heating  of  liquids  in  vessels  of 
glass  or  porcelain.  Eeduce  the  size  of  the  flame  as  the 
liquid  becomes  less. 

When  only  one-fourth  of  the  liquid  remains,  let  it  cool. 
If  nothing  happens  on  cooling,  evaporate  carefully  one- 
half  of  the  remaining  liquid  and  allow  to  cool  again. 
Describe  in  your  notes  the  crystals  which  separate.  They 
are  called  zinc  sulphate,  and  they  contain  zinc,  sulphur, 
and  oxygen.  Sulphuric  acid  contains  hydrogen,  sulphur, 
and  oxygen,  so  that  we  may  describe  the  change  which 
takes  place  in  the  gas-generating  bottle  by  the  statement 
that  the  zinc  takes  the  place  of  the  hydrogen. 


CHAPTER   VI 

OXYGEN  AND  HYDROGEN  PEROXIDE 

EXPERIMENT  23. — Preparation  and  properties  of  oxy- 
gen.— Place  about  a  gram  of  potassium  chlorate — free 
from  dirt — in  a  clean,  dry  test-tube  and  apply  a  gentle 
heat.  At  first  the  tube  must  be  simply  brushed  with  the 
flame.  Notice  the  decrepitation  and  melting  of  the  salt. 
Apply  the  spark  test.  (?)  When  no  more  gas  escapes  ex- 
amine the  residue  in  the  tube.  It  is  potassium  chloride. 
Compare  its  taste  with  that  of  potassium  chlorate.  Heat 
some  of  it  in  a  small  tube  of  hard  glass,  sealed  at  one  end. 
Does  it  give  off  oxygen  ?  Does  it  melt  as  readily  as  potas- 
sium chlorate? 

In  a  dry,  clean  test-tube  melt  carefully  about  a  gram 
of  potassium  chlorate,  remove  from  the  flame  and  show 
that  oxygen  is  not  escaping  by  the  spark.  Now  throw  into 


34:  ELEMENTARY  CHEMISTRY 

the  tube  about  ^  gram  of  powdered  manganese  dioxide 
and  test  again  with  the  spark.  (?)  What  is  catalytic  action? 

Make  a  mixture  of  two  parts  of  potassium  chlorate 
and  one  part  of  manganese  dioxide,  by  weight.  Both  sub- 
stances should  be  coarsely  powdered  and  free  from  dirt. 
Fill  a  wide  test-tube  to  the  depth  of  about  6  cm.  with  this 
mixture,  make  a  channel  by  tapping  on  the  table,  and 
clamp  the  tube  horizontally  near  the  mouth.  Provide  a 
well-fitting  cork  with  a  delivery  tube  to  collect  the  gas 
over  water.  Brush  the  tube  with  the  flame.  Rapid  heat- 
ing or  the  presence  of  dirt  in  materials  or  tube  may  cause 
dangerous  explosions.  When  the  gas  begins  to  escape  rap- 
idly, remove  the  flame,  and  do  not  begin  heating  again 
until  the  evolution  of  gas  slackens  or  ceases.  Only  a  very 
moderate  heat  is  required.  A  high  temperature  will  melt 
the  tube  and  spoil  the  experiment. 

Collect  the  gas  in  wide-mouthed  bottles.  Let  them 
remain  with  mouths  under  water  until  wanted,  or,  if  there 
is  not  room  for  this,  cover  them  with  wet  filter-paper  plas- 
tered tightly  over  the  mouth.  Also  collect  half  a  test-tube 
full  of  the  gas  to  investigate  its  solubility  in  the  same  way 
as  with  hydrogen.  If  the  oxygen  supply  fails,  remove  the 
exhausted  material  in  the  test-tube  by  poking  it  with  the 
handle  of  a  spoon,  put  in  fresh  material  and  again  heat 
gently.  Record  the  properties  of  the  gas,  including  its 
color,  odor,  and  solubility.  The  cloudiness  which  it  shows 
at  first  is  due  to  dust  carried  over  from  the  test-tube  and 
will  disappear  on  standing.  Inhale  a  little  of  the  gas. 

Heat  a  piece  of  charcoal  in  a  deflagrating  spoon  until 
it  glows  at  one  point.  Remove  it  from  the  flame  and 
plunge  it  into  a  bottle  of  oxygen.  (?)  Cover  the  jar  with 
a  glass  plate  and  put  it  aside.  Take  another  piece  of 
charcoal  and  burn  it  in  a  bottle  of  air  in  a  similar  way. 
It  may  be  necessary  to  heat  the  charcoal  several  times. 
Cover  this  bottle  also.  Examine  both  bottles.  What  is 


OXYGEN  AND  HYDROGEN  PEROXIDE  35 

evidently  the  nature  of  the  product  of  the  burning  of  char- 
coal ?  Into  both  bottles  pour  a  little  lime-water  and  shake. 
(  ?)  Cloudiness  produced  in  lime-water  is  evidence  of  the 
presence  of  carbon  dioxide  (so-called  carbonic-acid  gas). 
Does  charcoal  produce  the  same  substance  when  it  burns 
in  oxygen  as  when  it  burns  in  air?  Does  it  produce  the 
same  amount  of  heat  ?  Why  is  the  burning  in  oxygen  more 
rapid  and  brilliant? 

Put  some  sulphur  in  your  spoon  and  start  it  burning 
in  the  air.  Place  the  burning  sulphur  in  a  covered  bottle 
containing  air  and  allow  it  to  burn  for  a  time.  Cover  the 
jar  and  stand  it  aside.  Place  the  burning  sulphur  in  a 
bottle  of  oxygen.  (?)  Compare  the  two  bottles.  What 
is  the  nature  of  the  combustion-product  of  sulphur?  No- 
tice carefully  the  odor  of  the  contents  of  each  bottle. 
The  odor  is  extremely  irritating  and  suffocating.  Shake 
up  a  little  water  in  each,  and  add  to  the  water  a  small 
piece  of  red  and  a  small  piece  of  blue  litmus  paper.  (?) 
Heat  the  spoon  red-hot  to  remove  the  sulphur  and  allow  it 
to  cool  perfectly  before  using  it  again  or  putting  it  away. 

Straighten  a  watch-spring  by  pulling  it  out  into  a  line 
and  drawing  it  slowly  through  the  Bunsen  flame.  Slip 
aside  the  cover  of  a  bottle  of  oxygen  an  instant  and  throw 
in  a  layer  of  sand  1  cm.  deep.  Cut  a  piece  of  match-stick 
1  cm.  long  and  split  it  for  half  its  length.  Slip  this  over 
the  end  of  the  watch-spring,  set  fire  to  it,  and  introduce 
it  into  a  bottle  of  oxygen.  The  bottle  must  be  kept  covered 
during  the  experiment  and  the  spring  introduced  between 
bottle  and  cover.  If  nothing  happens,  place  another  split 
match-stick  on  the  spring  and  try  again.  What  is  the  re- 
sult? Describe  the  iron  oxide  produced. 

Put  some  magnesium  powder  in  a  small,  clean,  cold 
deflagrating  spoon.  The  powder  should  project  a  little 
above  the  edge  of  the  spoon.  Start  its  combustion  by  let- 
ting the  burner  flame  play  upon  it  and  let  it  burn  in  the 
25 


36  ELEMENTARY  CHEMISTRY 

air.  Let  cool,  and  examine  the  magnesium  oxide  produced. 
It  is  dangerous  to  remove  the  substance  until  perfectly 
cold.  Place  some  of  it  on  a  small  piece  of  red  litmus 
paper  and  wet  it.  After  a  minute  remove  the  magnesium 
oxide.  (?)  'Now  fill  the  spoon  again  with  magnesium 
powder,  start  it  burning,  and  plunge  it  into  a  bottle  of 
oxygen.  (?)  Examine  this  magnesium  oxide  and  test  it 
with  red  litmus  paper  as  you  did  the  other.  Are  the  two 
identical  ? 

Place  a  clean  piece  of  sodium  in  your  spoon — which 
must  be  cold  and  clean.  Heat  until  the  combustion  starts 
and  place  in  a  bottle  of  oxygen,  which  must  be  kept  cov- 
ered during  the  experiment.  After  the  combustion,  intro- 
duce a  little  water  into  the  bottle  and  shake.  Test  the 
liquid  with  red  and  blue  litmus  paper.  Use  very  small 
pieces. 

Sulphur  is  a  non-metal.  Notice  that  its  oxide  com- 
bines with  water  to  produce  a  substance  which  turns  lit- 
mus red — a  substance  which  is  an  acid.  Magnesium  and 
sodium  are  metals.  Their  oxides  combine  with  water  to 
produce  bases — substances  which  turn  litmus  blue.  This 
is  an  important  difference  between  metals  and  non-metals. 

EXPERIMENT  24. — Action  of  oxygen  upon  magnesium 
and  iron. — Cut  a  piece  of  asbestos  board  about  10  cm. 
square,  and  heat  it  for  five  minutes,  to  drive  off  moisture. 
Let  it  cool,  place  it  upon  the  trip-scales  or  a  lecture-table 
balance  and  place  upon  it  a  heap  of  magnesium  powder 
about  2  cm.  in  diameter.  Balance  the  other  pan  of  the 
scales  with  small  shot,  copper  filings  or  iron  filings.  Apply 
a  burner  flame  to  the  tip  of  the  heap.  (?)  Let  cool  and 
add  shot  or  filings  as  required  to  restore  equilibrium. 
Explain  exactly  what  has  occurred.  Repeat  the  same  ex- 
periment, using  powdered  iron  in  place  of  magnesium. 

EXPERIMENT  25. — Quantitative  experiment.  Decom- 
position of  potassium  chlorate.  Weight  of  a  liter  of  oxy- 


OXYGEN  AND   HYDROGEN   PEROXIDE 


3T 


PIG.  19. 


gen. — Before  the  student  can  make  the  calculations  con- 
nected with  the  following  experiment  he  must  understand 
the  methods  of  correcting  the  volumes  of  gases  for  tem- 
perature, pressure,  and  the  presence  of  water  vapor.  Un- 
less he  has  already  taken  up  this  subject  in  his  course  in 
physics,  he  should  study  the  discussion  in  the  Appendix 
and  solve  the  problems 
there  given.  Fit  up  the! 
apparatus  shown  in  Fig. 
19.  T  is  a  small  hard- 
glass  test-tube.  S  a  tight- 
ly fitting  stopper  of 
red  rubber,  the  tube 
from'  which  just  passes 
through  the  closely  fit- 
ting rubber  stopper  of 
the  acid  bottle  A.  This 
bottle  is  filled  with  water 
which  has  stood  in  the  laboratory  over-night,  and  is  ap- 
proximately at  room  temperature.  G  is  a  liter  graduated 
cylinder.  The  tube  A  A'  is  filled  with  water  by  placing  A 
in  a  beaker  of  water  and  .sucking  at  A'.  Then  the  Hoffmann 
clamp  C  is  tightened  and  the  stopper  R  tightly  inserted, 
the  tube  T  being  removed.  This  tube  is  cleaned  carefully, 
brushed  with  a  flame  until  warm,  and  accurately  weighed. 
Not  more  than  2.5  nor  less  than  1.5  grams  of  potassium 
chlorate  is  introduced  into  it,  and  it  is  brushed  again 
with  the  flame  until  the  potassium  chlorate  is  just  melted. 
The  heating  must  be  stopped  before  any  oxygen  escapes. 
Allow  the  tube  to  cool,  and  if  any  water  has  appeared  in 
the  upper  portion,  absorb  it  with  a  wisp  of  filter  paper. 
Weigh  the  tube  accurately.  Find  the  quantity  of  potas- 
sium chlorate  taken  by  subtraction. 

Place  the  cold  tube  tightly  on  the  stopper  S,  loosen  the 
clamp  C  and  cautiously  heat  the  potassium  chlorate,  finally 


38  ELEMENTARY  CHEMISTRY 

to  faint  redness.  The  secret  of  success  is  gradual  heating. 
Cool  gradually,  keeping  the  clamp  C  open  and  the  rubber 
tube  A'  below  the  level  of  the  water  in  G.  Disconnect  the 
apparatus.  Take  the  temperature  of  the  water  in  the  bottle 
A,  which  is  the  same  as  that  of  the  oxygen  collected  above 
it.  Read  the  volume  of  the  water  in  G,  which  is  that  of  the 
oxygen  (estimate  to  tenths  of  the  divisions  on  the  cylinder 
in  doing  this).  Eead  the  barometer  in  the  laboratory, 
which  will  give  you  the  atmospheric  pressure  under  which 
the  gas  was  collected.  Repeat  all  readings  to  be  sure  there 
is  no  mistake.  Record  all  readings  at  once  in  your  note- 
book. 

Weigh  the  tube  T  accurately  and  ascertain  by  subtrac- 
tion the  weight  of  the  oxygen  escaped  from  it.  Also  find 
by  subtraction  the  weight  of  potassium  chloride  it  con- 
tains. 

CALCULATION 

a.  How  much  oxygen  would  1  gram  of  potassium  chlo- 
rate yield? 

Solved  thus : 

Weight  oxygen  escaped          __ 
Weight  potassium  chlorate  taken  ~ 

Has  potassium  chlorate  always  the  same  composition? 
Does  your  result  agree  fairly  with  those  of  others  who 
have  made  the  same  experiment?  Why  not  exactly? 

b.  What  is  the  weight  of  1  liter  of  oxygen  just  as  you 
have  collected  it — that  is,  saturated  with  water  and  at  the 
temperature  and  pressure  of  the  air  of  the  laboratory? 

Solved  thus: 

Weight  oxygen  escaped  _ 
Volume  oxygen  in  liters  ~~ 

c.  What  is  the  weight  of  a  liter  of  oxygen  under  stand- 
ard conditions — that  is  dry,  at  a  temperature  of  0°  and  a 
pressure  of  760  mm.  of  mercury? 


OXYGEN  AND  HYDROGEN  PEROXIDE  39 

Solved  in  the  same  way  as  c,  but  you  must  first  find 
what  the  volume  of  your  oxygen  would  be  if  dry  and  at  0° 
and  760  mm.  This  you  can  do,  as  explained  in  the  Ap- 
pendix (p.  150),  by  the  formula: 

FX273XP-W 
(273  +  t)  X  760 

V0  =  standard  volume. 

t  =  observed  temperature  at  which  your  gas  was  meas- 
ured. 

P  =  observed  pressure. 

w  =  pressure  of  water  vapor  at  the  temperature  t  (this 
can  be  obtained  from  table  in  Appendix). 

The  following  will  serve  as  an  example  of  the  mode  of 
calculating  the  results : 

Volume  of  water  collected  in  G 600  c.c.,  or  .6  liter. 

Temperature 14°  C. 

Barometer 756  mm. 

The  weight  of  the  test-tube  containing  potas- 
sium chlorate,  dried  by  being  melted,  was  12.50  grams. 
Weight  of  empty  test-tube 10.46 

Potassium  chlorate  taken 2.04  grams. 

After    being   heated    to  redness   the  tube 
weighed 11.70  grams. 

Hence  the  weight  of  oxygen  which  escaped 

was . .   12.50  grams. 

11.70  grams. 
Weight  of  escaped  oxygen  =       .80  gram. 

a.  How  much  oxygen  would  1  gram  of  potassium 
chlorate  yield  ?  Clearly  the  answer  is  equal  to 

Weight  of  escaped  oxygen          _  .80  __ 
Weight  of  potassium  chlorate  taken  ~~  2.04  ~ 


40  ELEMENTARY  CHEMISTRY 

b.  What  is  the  weight  of  1  liter  of  oxygen  just  as  col- 
lected ?     The  answer  is  given  by 

Weight  of  oxygen  _. 80 gram  _ 
Volume  of  oxygen"  .6  liter-  ~ 

c.  What  is   the  weight  of   1   liter  of  oxygen  under 
standard  conditions — that  is,  at  0°  C.,  a  pressure  of  760 
mm.,  and  in  perfectly  dry  condition  ? 

First  find  the  volume  of  your  oxygen  at  standard  con- 
ditions.   Use  the  formula 


0  ~      (273  +  t)  X  760) 
F=*6  liter^ - 
t  =  14°. 
P  =  756  mm. 

w  (the  pressure  of  water  vapor  at  14°)  we  find  from 
table  in  Appendix  =  11.9  mm. 

Substituting  these  values, 

J7      600  X  273  X  (756  -  11.9) 

287  X  760  =          CiC*'  °r 

The  weight  of  1  liter  of  dry  oxygen  at  0°  and  760  mm. 
would  be 

.80  gram  _ 
.559  liter  ~ 

Do  not  expect  your  results  to  be  identical  with  those 
given  above.  The  problem  is  simply  to  illustrate  the 
method  of  calculation,  and  the  numbers  were  chosen  arbi- 
trarily to  give  results  quite  close  to  the  truth. 

Very  careful  work  has  shown  that 
the  weight  of  a  liter  of  dry  oxy- 
gen at  0°  and  760  mm.  is 1.429  grams. 

And  that  1  gram  of  pure  potassium 
chlorate  gives  off  when  heated . .  .392  gram  of  oxygen. 


OXYGEN  AND   HYDROGEN   PEROXIDE  41 

Your  results  should  not  vary  widely  from  these  figures. 
If  they  do  the  experiment  must  be  repeated;  but  before 
repeating  it  go  over  the  calculation  with  great  care,  for  the 
error  is  more  likely  to  be  found  there  than  in  the  weights 
or  measurements. 

EXPERIMENT  26. — Changes  produced  in  air  by  respira- 
tion.— Place  some  clean  lime-water  in  a  beaker,  blow  gently 
through  it  by  means  of  a  glass  tube.  (?)  Invert  three 
wide-mouthed  bottles  full  of  water  in  your  tin  pan  and 
collect  air  in  the  first  bottle  from  the  beginning  of  an 
expiration.  Withdraw  the  bottle  from  the  water  by  means 
of  a  glass  plate,  slip  the  plate  aside  an  instant  and  lower 
into  the  bottle  a  lighted  candle.  Does  it  ta*n  as  long  as 
it  would  in  the  same  volume  of  pure  air?  Why? 

In  the  second  bottle  collect  air  from  the  end  of  an 
expiration,  using  the  last  portions  of  air  from  the  lungs. 
Test  this  with  a  candle.  ( ?)  The  result  is  due  partly  to 
the  small  amount  of  oxygen,  and  partly  to  the  large  amount 
of  carbon  dioxide. 

Fill  the  lungs  with  air  and  hold  the  breath  as  long  as 
you  can  without  discomfort.  Collect  the  first  of  the 
expiration  and  test  it  with  a  candle.  Does  the  result  fur- 
nish any  evidence  of  the  diffusion  of  carbon  dioxide  up- 
ward, or  of  oxygen  downward,  in  the  lungs? 

The  body  burns  up  about  220  grams  of  carbon  in 
twenty-four  hours,  almost  all  of  which  is  cast  out  through 
the  lungs  as  carbon  dioxide.  Weigh  off  roughly  this 
amount  of  charcoal  on  the  trip-scales  in  order  to  get  an 
idea  of  the  quantity.  Of  course  the  amount  varies  greatly 
in  different  people  and  in  the  same  person  at  different 
times.  The  greater  the  activity  and  the  lower  the  tem- 
perature'of  the  surrounding  air,  the  more  active  the  in- 
ternal combustion  becomes. 

EXPERIMENT  27. — Ozone. — Set  up  the  apparatus  illus- 
trated in  Fig.  16,  Part  I,  and  perform  the  experiment 


42  ELEMENTARY  CHEMISTRY 

described  in  the  first  part  of  paragraph  54.  Make  a  sketch 
of  your  apparatus  and  write  a  description  of  your  work. 
The  oxygen  required  must  be  obtained  from  a  cylinder  of 
the  compressed  gas,  not  made  from  potassium  chlorate. 
If  there  is  no  cylinder  available,  use  air  instead,  but  the 
proof  is  not  conclusive.  The  induction  coil  can  be  excited 
by  two  dichromate  or  three  Edison-Lalande  cells,  or  in 
any  other  suitable  way.  Before  beginning  the  experiment 
read  the  directions  carefully,  make  a  list  of  the  things  you 
need  and  procure  them,  so  as  not  to  be  interrupted  at  a 
critical  point  by  the  lack  of  something  essential.  The 
solution  of  starch  and  potassium  iodide  must  be  very 
dilute — not  more  than  1  gram  of  starch  and  .2  gram  of 
potassium  iodide  in  300  c.c.  of  water.  It  should  be  per- 
fectly clear  and  limpid. 

Procure  an  ozone  tube  (Fig.  16,  Part  I)  and  prepare 
ozone  as  described  in  paragraph  54,  Part  I.  The  apparatus 
is  shown  in  Fig.  17,  Part  I.  The  current  of  oxygen  comes 
from  a  cylinder  or  a  gasometer  and  must  be  the  gentlest 
possible.  The  slower  the  current  the  richer  the  gas  will 
be  in  ozone,  since  it  spends  a  longer  time  under  the  in- 
fluence of  the  discharge.  The  preparation  of  pure  ozone 
by  intense  cooling  is  too  dangerous  and  difficult  to  at- 
tempt. Instead  of  this,  note  carefully  the  odor  of  the 
issuing  gas  and  hold  a  piece  of  starch  potassium  iodide 
paper  in  it. 

Place  barium  dioxide  to  the  depth  of  about  J  cm.  in  a 
test-tube  and  allow  some  strong  sulphuric  acid  to  trickle 
down  the  side  of  the  tube.  Test  the  gas  given  off  with 
starch  potassium  iodide  paper  and  notice  its  odor.  Cover 
the  tube  with  a  clean  silver  coin  and  stand  it  in  a  rack 
until  the  end  of  the  laboratory  period.  If  there  is  no  ap- 
parent result,  do  not  attribute  the  fact  to  an  error  in  your 
work.  The  action  requires  a  long  time  and  a  gas  rich  in 
ozone.  Sketch  the  apparatus  in  your  note-book. 


OXYGEN  AND  HYDROGEN  PEROXIDE  43 

What  experimental  evidence  have  you  that  ozone  and 
oxygen  are  two  forms  of  the  same  element?  Which  is  the 
stable  form  of  oxygen  (1)  at  ordinary  temperatures?  (2) 
at  slightly  elevated  temperatures  (say  300°)  ?  (3)  at  very 
high  temperatures? 

The  existence  of  the  same  element  in  two  or  more  dif- 
ferent modifications  is  called  allotropy.  Do  you  know  of 
any  other  elements  which  exist  in  allotropic  modifications? 
Can  you  think  of  any  allotropic  modifications  of  carbon? 
Of  tin?  When  the  substance  is  a  compound  the  existence 
of  several  modifications  is  called  isomerism.  What  sub- 
stance have  we  studied  which  exists  in  isomeric  modifica- 
tions ? 

Rub  the  head  of  a  match  gently  with  the  moistened 
finger.  Notice  the  odor.  (?)  Consult  the  instructor  about 
the  preparation  of  ozone  by  the  action  of  phosphorus  on 
the  oxygen  of  the  air. 

EXPERIMENT  28. — Hydrogen  peroxide. — Examine  com- 
mercial hydrogen  peroxide.  Pour  a  little  of  it  upon  some 
manganese  dioxide  in  a  test-tube  and  test  the  gas  given  off 
with  a  spark.  (?)  Take  about  25  c.c.  of  hydrogen  perox- 
ide in  a  beaker  and  dissolve  in  it  a  piece  of  sodium  hy- 
droxide about  1  cm.  in  length.  This  is  to  neutralize  the 
acid  which  hydrogen  peroxide  always  contains,  and  which 
would  interfere  with  the  experiments.  Add  some  of  this 
liquid  to  some  water  weakly  colored  with  aniline  red.  (?) 
Be  sure  that  the  water  does  not  contain  any  solid  aniline 
red  at  the  bottom,  for  if  it  does,  the  solid  will  continually 
dissolve  and  spoil  the  result. 

Put  the  rest  of  your  hydrogen  peroxide  into  a  wide  test- 
tube  provided  with  a  cork  and  a  delivery  tube  leading  to  a 
vessel  full  of  water.  Boil  gently  and  collect  the  gas  in 
inverted  test-tubes  filled  with  water.  Test  the  gas  with 
the  spark.  (?)  Only  the  hydrogen  peroxide  takes  part  in 
the  change.  What  has  happened?  What  remains  in  the 


44  ELEMENTARY  CHEMISTRY 

test-tube  in  which  the  hydrogen  peroxide  was  boiled? 
Why  is  the  cork  frequently  driven  out  of  the  stock-bottle 
of  hydrogen  peroxide  ? 

State  precisely  the  relation  in  composition  between 
hydrogen  peroxide  and  Water.  State  precisely  what  is 
meant  by  catalytic  action,  using  hydrogen  peroxide  as  an 
illustration. 

PKOBLEMS 

10.  If  75  c.c.  of  oxygen  could  be  transformed  completely  into 
ozone,  what  volume  of  ozone  would  be  obtained  ? 

11.  115  c.c.  of  oxygen  were  partly  converted  into  ozone.     The 
volume  contracted  to  110  c.c.,  but  when  the  gas  was  gently  heated 
by  a  burner  flame  the  original  volume  was  exactly  restored.     Cal- 
culate («)  the  volume  of  ozone  which  had  been  produced  and  (&) 
its  percentage  by  volume  in  the  gas. 

12.  160  c.c.  of  oxygen  containing  ozone  were  heated.     The 
volume  became  170  c.c.     (a)  How  many  c.c.  of  ozone  and  (&)  what 
percentage  of  it  by  volume  were  present  ? 


CHAPTER   VII 

COMBUSTION 

EXPERIMENT  29. — Combustion. — Fit  up  a  lamp  chim- 
ney with  a  perforated  rubber  cork  and  tube  for  supplying 
gas  (Fig.  19,  p.  52,  Part  I).  Cover  the  chimney  with  a 
piece  of  asbestos  having  a  hole  in  it  large  enough  to  admit 
your  spoon.  Clamp  the  chimney  at  its  upper  portion  and 
see  that  the  cork  is  tightly  placed  in  the  bottom.  Use  illu- 
minating gas — not  hydrogen — and  make  the  experiments 
on  combustion  described  in  Chapter  VII,  paragraphs  02 
nn<l  03,  Part  I. 

In  the  experiment  with  nitric  acid  use  fuming  nitric 
acid.  The  ordinary  acid  is  not  strong  enough  for  the  pur- 


COMBUSTION  45 

pose.  Keep  the  burning  substances  out  of  contact  with 
the  glass  to  avoid  breakage.  Always  hold  your  spoon  by 
the  upper  end,  and  keep  the  hand  some  distance  above 
the  chimney,  to  avoid  burning  the  fingers. 

In  getting  the  oxygen  flame  (Fig.  20,  Part  I)  use  a 
mouth  blowpipe  the  tip  of  which  has  been  carefully  bent 
until  it  is  parallel  with  the  longer  portion.  Connect  this 
with  an  oxygen  cylinder  and  obtain  the  gentlest  current 
of  oxygen  possible.  Test  the  strength  of  the  current  by 
letting  it  blow  on  the  moistened  finger.  Insert  the  blow- 
pipe through  the  flame  into  the  chimney.  Keep  the  oxy- 
gen flame  away  from  the  glass. 

The  same  chimney  can  be  used  for  the  air-flame;  but 
you  will  need  a  doubly  perforated  cork,  one  end  of  which 
carries  a  wide,  straight  tube  open  at  both  ends  (Fig.  21, 
p.  55,  Parti). 

Hold  a  small  crystal  of  each  of  the  four  chlorates 
mentioned  in  paragraph  63,  Part  I,  in  the  lower  part 
of  the  Bunsen  flame  with  forceps.  Kecord  the  flame  colors 
of  the  corresponding  metals. 

What  is  a  flame?  Why  do  some  substances  burn  with 
flame  and  others  not?  Explain  what  is  meant  by  the 
statement  that  the  combustion  of  one  gas  in  another  is  re- 
versible. Explain  exactly  what  is  meant  by  the  statement 
that  the  air  flame  is  the  gas  flame  turned  inside  out. 

EXPERIMENT  30. — The  slow  combustion  of  iron. — Stuff 
a  wad  of  steel  wool  about  half-way  up  a  tube  closed  at  one 
end  and  graduated.  Wet  the  steel  and  clamp  the  tube, 
open  end  down,  in  a  vessel  of  water.  The  level  of  the 
water  inside  the  tube  must  be  within  the  graduations. 
If  it  is  not,  remove  the  tube  from  the  vessel,  pour  some 
more  water  in  the  tube,  cover  the  end  with  the  thumb,  put 
it  again  in  the  vessel  of  water  and  clamp  it.  Read  the 
level  of  the  water  inside  the  tube,  allow  to  stand  for  half 
an  hour  or  as  long  a  time  as  possible,  and  read  the  level 


46  ELEMENTARY  CHEMISTRY 

again.  Examine  the  steel  carefully.  What  does  the  result 
show  ?  Predict  the  composition  of  rust.  Was  heat  evolved 
or  absorbed  in  this  experiment  ?  Steel  is  chiefly  composed 
of  iron,  and  it  was  the  iron  alone  which  was  affected  in 
the  experiment.  Dislodge  the  steel  wool  with  the  handle 
of  a  spoon,  or  with  a  glass  rod,  and  clean  the  tube  with 
a  small  quantity  of  strong  hydrochloric  acid. 


CHAPTER   VIII 


No  experiments. 


PROBLEMS 


Calculate  the  percentage  composition  of  the  substances  whose 
formulas  are  given  below.  The  amount  of  each  constituent  must 
be  obtained  by  an  independent  calculation,  never  by  subtracting 
the  sum  of  the  others  from  100.  State  the  results  to  two  decimal 
places.  If  the  third  decimal  place  is  greater  than  5  add  1  to  the 
second ;  if  less  than  5  discard  it. 

13.  Mercuric  oxide HgO 

14.  Water H,O 

15.  Mercuric  chloride HgCU 

16.  Mercurous  chloride HgCl 

17.  Potassium  chlorate KC1O8 

18.  Table  salt  (sodium  chloride) NaCl 

19.  Manganese  dioxide MnOs 

Calculate  the  simplest  formulas  of  the  following  substances  from 
their  percentage  compositions. 

44.07  per  cent. 

55.93  " 

40  " 

12  " 

48  " 

52.35  " 

47.65  " 


20.  Hydrogen.. 
Chlorine  .  .  . 
21.  Nitrogen.  .  . 
Oxygen  
22.  Carbon... 

2.74  per  cent. 
97.26        " 
30.43        " 
69.57        " 
40.00        " 

23.  Mercury.. 
Iodine,  .  .  . 
24.  Calcium.. 
Carbon.  .  . 

Hydrogen  .  . 
Oxygen  

6.67        " 
53.33        « 

Oxygen  .  . 
25.  Potassium 
Chlorine.  . 

SALT  AND  SODIUM 


CHAPTER    IX 

SALT  AND   SODIUM 

Remember  the  precautions  necessary  in  using  sodium 
(p.  28).  Carry  out  this  experiment  under  the  hood. 
Avoid  inhaling  chlorine. 

EXPERIMENT  31. — Salt  and  sodium. — Fit  up  the  appa- 
ratus shown  in  Fig.  20,  and  carry  out  the  synthesis  of  salt. 
The  flask  should  be  small.  It  is  one-fourth  filled  with 
manganese  dioxide  in  small  lumps — not  powder.  This 
must  be  allowed  to  slide  slowly  into  the  flask,  the  latter 
being  in  an  inclined  position.  The  perforated  cork  and 
exit  tube  must  fit  tightly;  otherwise  chlorine  will  escape 
and  cause  great  distress 
and  possibly  injury. 
The  bulb  must  be  of 
hard  glass.  It  con- 
tains a  fragment  of  so- 
dium the  size  of  a  pea, 
which  may  be  elongated 
by  rolling,  to  get  it  into 
the  tube. 

Cover  the  manga- 
nese dioxide  with  strong 
hydrochloric  acid,  place 
the  flask  on  wire  gauze,  connect  the  apparatus,  and  apply  a 
gentle  heat.  The  tube  which  is  bent  downward  dips  into 
about  20  c.c.  of  alcohol  or  sodium  hydroxide  solution, 
which  will  absorb  the  excess  of  chlorine. 

When  the  gas  in  the  bulb  appears  green  (?)  heat  the 
sodium  very  gently  with  a  small  flame.  When  the  action 
begins,  stop  heating  the  sodium.  When  the  action  is  over, 
extinguish  the  burner,  take  out  the  cork  of  the  chlorine 
generator  and  fill  the  latter  with  water  to  stop  the  evolu- 


FIG.  20. 


48  ELEMENTARY  CHEMISTRY 

tion  of  gas.  It  is  well  to  do  this  quickly,  to  hold  the 
breath  during  the  operation,  and  to  retire  ten  feet  or  more 
for  a  short  time  until  the  chlorine  which  the  water  forces 
out  of  the  flask  is  dissipated. 

Break  open  the  bulb.  The  black  product  is  silicon 
due  to  the  unavoidable  action  of  the  sodium  upon  the 
glass.  Taste  the  white  product.  Is  it  soluble  in  water  ?  Is 
it  possible  for  an  insoluble  substance  to  have  any  taste? 
Why?  What  information  does  this  experiment  give  re- 
garding the  composition  of  salt  ?  Is  the  information  qual- 
itative or  quantitative  ?  Throw  the  fragments  of  the  bulb- 
tube  into  a  vessel  of  water  to  get  rid  of  any  sodium  which 
may  have  escaped  the  action  of  the  chlorine. 

Does  salt  contain  more  or  less  energy  than  its  con- 
stituents taken  separately?  How  do  you  know?  If  you 
had  to  decompose  salt  into  sodium  and  chlorine,  would  it 
be  necessary  to  supply  energy  or  would  the  decomposition 
furnish  energy  which  could  be  applied  to  other  purposes? 
Recall  all  the  cases  in  your  past  work  in  which  you  have 
brought  about  or  witnessed  the  combination  of  two  ele- 
ments. Was  energy  absorbed  or  given  out  in  these  ex- 
periments ?  Recall  the  cases  in  which  a  compound  has  been 
decomposed.  Was  energy  given  out  or  absorbed?  Irrita- 
tion from  chlorine  can  be  relieved  by  at, once  inhaling  the 
vapor  of  alcohol. 

EXPERIMENT  32. — Preparation  of  pure  salt. — Prepare 
a  saturated  solution  of  salt,  filter  the  liquid,  and  pass 
hydrochloric  acid-gas  into  it  through  an  inverted  funnel, 
whose  rim  just  dips  into  the  liquid.  The. gas  is  obtained 
as  described  on  p.  56,  Experiment  41,  the  sulphuric  acid 
being  allowed  to  drop  very  slowly  into  the  hydrochloric. 
Salt,  being  insoluble  in  hydrochloric  acid,  is  precipitated; 
the  other  substances  present  remain  dissolved.  When 
enough  salt  has  collected,  pour  away  the  liquid  and  wash 
the  salt  three  times  with  very  small  quantities  of 


SALT  AND  SODIUM  49 

tilled  water.  (Why  is  this  necessary?)  Place  the  salt  in 
a  clean  dish  and  apply  heat  with  a  small  flame  which  does 
not  touch  the  dish,  stirring  the  salt  constantly  with  a 
glass  rod.  Preserve  the  pure  salt  in  a  dry,  clean  test- 
tube,  tightly  corked  and  labelled. 

EXPERIMENT  33.  —  Sodium.  —  Examine  sodium,  and 
write  a  description  of  it.  Cut  it.  Is  it  a  metal?  Why 
does  the  luster  fade  ?  How  could  the  fading  be  prevented  ? 
Could  sodium  be  used  in  the  same  way  as  iron  or  copper  ? 
Why  ?  How  does  its  density  compare  with  that  of  familiar 
metals  ? 

Hold  a  small  fragment  of  sodium  in  the  flame  in  for- 
ceps. (?)  Explain,  from  previous  work,  the  behavior  of 
sodium  with  oxygen,  with  chlorine,  and  with  water.  Ex- 
plain precisely  what  is  meant,  both  qualitatively  and  quan- 
titatively by  the  equation— 


EXPERIMENT  34.  —  Sodium  compounds.  —  Sodium  hy- 
droxide must  not  be  touched  with  the  fingers.  The  bottle 
containing  it  "must  be  kept  tightly  corked.  Examine  so- 
dium hydroxide,  sodium  nitrate,  sodium  carbonate  (anhy- 
drous), sodium  carbonate  (crystallized),  and  sodium  sul- 
phate. Are  they  soluble  in  water?  What  color  does  each 
communicate  to  the  flame?  Use  a  clean  iron  wire  hot 
enough  to  make  a  fragment  of  the  sodium  compound  ad- 
here to  it.  After  the  test,  clean  the  wire  by  dipping  it 
into  a  little  hydrochloric  acid  in  a  test-tube  (not  in  the 
bottle)  and  holding  it  in  the  flame  until  the  latter  shows 
no  color,  and  use  it  to  test  the  next  substance. 

Prepare  a  solution  of  sodium  hydroxide  by  dissolving 
a  piece  about  1  cm.  long  in  50  c.c.  of  water  in  a  beaker. 
Test  small  separate  portions  of  it  with  blue  litmus  paper, 
with  red  litmus  paper  (use  small  pieces),  with  a  drop  of 
phenol  phthalein  solution,  with  a  drop  of  cochineal  solu- 


50  ELEMENTARY  CHEMISTRY 

tion.  What  are  the  results?  Liquids  which  affect  these 
coloring  matters  in  this  way  are  called  alkaline. 

Place  the  rest  of  the  sodium  hydroxide  solution  in  a 
clean  dish  and  add  hydrochloric  acid  to  it  one  drop  at  a 
time.  After  each  drop  stir  with  a  glass  rod  and  test  the 
liquid  with  a  fragment  of  blue  litmus  paper.  When  the 
paper  is  just  turned  red,  remove  it  and  evaporate  the  liquid 
carefully  to  dryness.  Turn  down  the  flame  very  much  to- 
ward the  end  and  keep  it  moving.  Taste  the  residue  when 
it  is  perfectly  dry.  What  is  it?  What  other  product 
must  have  been  formed?  Write  the  equation. 

Expose  a  clean  crystal  of  sodium  carbonate  and  a  piece 
of  sodium  hydroxide  on  two  separate  watch  glasses  to  the 
air  for  an  hour  or  overnight.  What  happens?  What  is 
meant  by  efflorescence?  by  deliquescence?  What  has  the 
state  of  the  air  to  do  with  both  ?  How  would  any  soluble 
salt  behave  if  kept  in  an  atmosphere  saturated  with  water- 
vapor  ? 

EXPERIMENT  35. — Sodium  carbonate.  Sodium  acid 
carbonate. — Stir  up  some  ammonia  water  with  salt  in  a 
beaker  until  the  solution  is  saturated.  Let  the  undissolved 
salt  settle  and  pour  off  the  solution  into  another  beaker. 
Pass  carbon  dioxide  into  it  through  a  funnel  the  rim  of 
which  dips  into  the  liquid.  The  carbon  dioxide  can  be  ob- 
tained from  a  Kipp  generator  or  made  in  a  gas-generating 
bottle  from  marble  and  dilute  hydrochloric  acid  as  directed 
on  p.  133,  Experiment  127.  The  precipitate  is  sodium 
acid  carbonate  (baking  soda),  NaHC03.  Filter  and  dry 
the  product  by  pressing  it  between  layers  of  filter  paper, 
which  are  renewed  as  they  become  wet. 

Arrange  a  test-tube  with  a  cork  and  a  delivery-tube 
dipping  into  a  little  lime-water.  Place  some  dry  sodium 
acid  carbonate  in  it,  support  the  tube  horizontally,  and 
apply  a  very  gentle  heat.  What  gas  is  given  off?  What 
condenses  in  the  tube?  Compare  the  residue  in  the  tube 


wi 

; 


SALT   AND  SODIUM  51 


VSi 

- 


with  anhydrous  sodium  carbonate,  Na2C03.     Write  the 

uation. 

These  two  reactions  illustrate  the  ammonia  soda  proc- 
ess, the  most  important  method  of  manufacturing  sodium 

rbonate. 

Make  a  mixture  of  sodium  acid  carbonate  with  about 
twice  its  weight  of  cream  of  tartar.  Grind  the  two  to- 
gether in  a  dry  mortar.  What  familiar  substance  have  you 
made?  Is  there  any  action?  As  a  rule,  chemical  action 
between  solids  at  ordinary  temperature  and  pressure  is 
slow,  so  slow  that  it  is  almost  imperceptible.  Place  some 
of  the  mixture  in  a  test-tube  and  add  water.  (?)  Fit  a 
cork  bearing  a  delivery  tube  to  the  test-tube  and  pass  some 
of  the  gas  into  lime-water.  (?)  What  is  the  function  of 
this  gas  in  the  baking  process  ?  The  same  gas  is  produced 
when  sodium  acid  carbonate  is  treated  with  an  acid — e.  g., 
hydrochloric  acid.  (Try  it.) 

EXPERIMENT  36. — Water  of  crystallization. — Examine 
some  crystallized  copper  sulphate.  Place  about  2  grams  of 
it  in  a  dry,  clean  dish,  cover  with  a  dry,  clean  beaker,  and 
heat  by  a  small  flame,  kept  in  motion.  (?)  After  about 
five  minutes,  remove  the  beaker  and  heat  the  substance  a 
little  more  intensely,  stirring  it  and  pressing  the  lumps 
gently  with  a  glass  rod.  Let  the  product  cool  and  compare 
it  with  the  original  substance.  What  is  it  ?  Sprinkle  some 
water  upon  it.  (?)  It  becomes  hot.  What  is  the  source 
of  the  heat?  What  is  water  of  crystallization  and  why  is 
it  so  called  ?  From  the  facts  furnished  by  this  experiment, 
discuss  the  question  whether  water  of  crystallization  is 
chemically  combined  or  merely  mixed. 

PROBLEMS 
26.  Calculate  the  percentage  composition  of 

a.  Sodium  carbonate NaaCO8, 

&.  Cystallized  sodium  carbonate NaaCOs,  10HaO. 

In  b  calculate  water,  not  hydrogen. 
26 


52  ELEMENTARY  CHEMISTRY 

27.  How  much  sodium  carbonate  can  be  made  from  10  kilos 
of  salt  ? 

28.  How  much  sodium  is  necessary  to  decompose  36  grams  of 
water,  and  how  much  hydrogen  is  liberated  ? 

29.  How  much  sodium  can  be  obtained  by  the  electrolysis  of 
20  grams  of  sodium  hydroxide  ? 

30.  A  piece  of  sodium  was  placed  in  water.    500  c.c.  of  hydro- 
gen under  standard  conditions  escaped.     What  was  the  weight  of 
the  sodium  ? 

31.  How  much  sodium  carbonate  and  slaked  lime  are  needed 
to  make  2  kilos  of  sodium  hydroxide  ? 


CHAPTER    X 

CHLORINE 

EXPERIMENT  37. — Chlorine. — Carry  out  this  experi- 
ment under  the  hood.  Avoid  inhaling  the  gas.  Be  sure 
that  the  cork  of  the  apparatus  fits  tightly.  If  the  throat 
becomes  irritated,  place  some  alcohol  in  a  beaker  or  on  a 
handkerchief  and  inhale  the  vapor. 

Place  crystals  of  potassium  permanganate  1  cm.  deep 
in  a  flask  (|  liter  capacity  or  less).  The  substance  need 
not  be  pure.  Fit  the  flask  with  a  rubber  stopper  carrying 
a  dropping  funnel  and  a  delivery  tube.  Support  it  firmly 
on  wire  gauze,  and  put  a  small  flame  8  cm.  below  it. 
Allow  hydrochloric  acid  to  fall  drop  by  drop  into  the  flask. 
Collect  by  downward  displacement  four  wide  test-tubes 
and  four  bottles  full  of  the  gas.  Judge  when  the  vessel  is 
full  by  the  color,  and  immediately  cover  it  tightly  with  a 
glass  plate  and  substitute  another;  otherwise  the  excess  of 
chlorine  will  be  forced  out  at  the  top  and  make  the  air 
unfit  to  breathe.  Test-tubes  full  of  chlorine  can  be  corked. 
The  exit  tube  of  the  chlorine  apparatus  must  reach  to  the 
bottom  of  the  vessel.  This  tube  should  be  cut  about  a  foot 


CHLORINE  53 

from  the  desk  and  united  again  by  a  short  piece  of  rubber 
tube,  that  it  may  be  moved  without  disturbing  the  appa- 
ratus. The  vessel  in  which  you  are  collecting  chlorine 
should  be  kept  covered,  and  the  exit  tube  slipped  between 
the  cover  and  the  side  of  the  bottle.  The  bottles  of  chlorine 
must  be  kept  covered  during  all  experiments.  When  you 
have  finished  and  desire  to  get  rid  of  the  chlorine,  place 
the  bottles  under  the  hood,  remove  the  covers  without 
breathing,  ami  at  once  retire  to  a  distance.  After  the  chlo- 
rine has  escaped — which  will  require  ten  minutes — the 
bottles  can  be  cleaned. 

Record  the  physical  properties  of  the  element.  Invert 
a  test-tube  full  in  water  and  shake  gently.  Is  it  soluble? 
Into  a  test-tube  filled  with  the  gas  throw  a  pinch  of  pow- 
dered arsenic.  ( ?)  Into  another  test-tube  a  little  pow- 
dered antimony.  (  ?)  Fill  a  small,  clean  test-tube  with 
hydrogen  from  the  Kipp  apparatus  by  upward  displace- 
ment. Without  losing  any  of  either  gas,  bring  this  tube 
mouth  to  mouth  with  a  test-tube  filled  with  chlorine.  The 
two  tubes  should  be  of  the  same  size,  and  the  tube  contain- 
ing the  hydrogen  uppermost.  Keeping  the  mouths  to- 
gether, invert  several  times  to  mix  the  gases,  then  hold 
the  mouth  of  each  tube  to  a  flame.  This  experiment  must 
not  be  made  in  direct  sunlight.  Why  not?  What  is  the 
product  ? 

Lower  a  hydrogen  flame  into  a  bottle  of  chlorine.  If 
you  use  a  gas-generating  bottle  in  this  experiment,  remem- 
ber that  it  must  be  allowed  to  run  five  minutes  and  that 
the  generator  and  stopper  must  be  wrapped  in  a  towel  be- 
fore lighting  the  hydrogen.  Describe  the  appearance  of 
the  flame  of  hydrogen  burning  in  chlorine.  Hold  a  glass 
rod  bearing  a  drop  of  ammonia  near  an  open  bottle  of 
hydrochloric  acid.  (?)  This  is  a  test  for  hydrochloric 
acid.  Now  hold  a  drop  of  ammonia  in  the  gas  left  in  the 
jar  in  which  the  hydrogen  was  burned.  Is  it  hydrochloric 


54  ELEMENTARY  CHEMISTRY 

acid  ?  Has  it  the  color  of  chlorine  ?  Heat  a  piece  of  char- 
coal in  a  spoon  until  it  begins  to  burn,  and  place  it  in 
chlorine.  (?) 

Take  out  the  charcoal  and  lower  a  lighted  candle  sup- 
ported on  a  wire  into  the  same  bottle  of  chlorine.  The 
candle  is  composed  chiefly  of  compounds  of  hydrogen  and 
carbon.  The  black  substance  thrown  off  from  the  flame 
is  carbon  (soot).  Use  the  two  preceding  results  to  explain 
this  one.  What  else  must  be  produced?  Let  the  covered 
jar  stand  till  the  soot  settles  and  try  the  test  with  the  glass 
rod  bearing  a  drop  of  ammonia. 

In  the  third  bottle  place  a  bit  of  red  litmus  paper,  a 
piece  of  blue  litmus  paper,  a  strip  of  colored  calico,  and  a 
fragment  of  printed  matter  smeared  over  with  writing  ink 
until  it  is  illegible.  Describe  and  explain  the  results.  What 
is  the  most  important  use  of  chlorine? 

In  the  fourth  bottle  place  some  fresh  slaked  lime  (milk 
of  lime  answers  well)  and  immediately  cover  the  bottle 
tightly  with  the  palm  of  the  hand  and  shake  it.  (?)  Does 
the  chlorine  disappear? 

Carefully  add  nitric  acid,  a  few  drops  at  a  time,  to  the 
contents  of  the  bottle.  What  happens?  What  very  im- 
portant technical  process  does  this  illustrate? 

EXPERIMENT  38. — Chlorine  water.  Bleaching. — Pre- 
pare some  chlorine  water  by  passing  chlorine  made  as  in 
Experiment  37  into  cold  water  under  the  hood  for  twenty 
minutes.  You  will  need  about  500  c.c.  Examine  it.  Has 
it  any  of  the  properties  of  the  gas  ?  Does  it  bleach  ?  Try 
small  portions  of  it  with  red  and  blue  litmus  paper.  With 
a  drop  of  ink. 

Select  a  tube  of  soft  glass  1  cm.  or  more  wide  and 
about  1  meter  long.  Seal  it  near  one  end  by  drawing  it 
out  in  the  flame  of  the  blast  lamp.  Cover  it  with  soot 
with  the  yellow  flame  before  cooling,  to  avoid  cracking. 
When  cool  wash  off  the  soot,  fill  the  tube  with  chlorine 


CHLORINE  55 

water  and  invert  it  in  the  same  liquid  in  a  glass  or  porce- 
lain dish.  Clamp  vertically  and  allow  to  stand  in  direct 
sunlight  as  long  as  possible. 

When  a  sufficient  quantity  of  gas  has  collected,  cover 
the  end  of  the  tube  with  the  thumb,  remove  it  from  the 
dish  and  test  the  gas  with  the  spark.  Explain  what  has 
happened. 

Prepare  a  thin  paste  of  bleaching  powder  and  water  in 
a  small  beaker  or  dish.  Add  a  few  drops  of  sulphuric 
acid  (why?)  and  soak  in  the  liquid  a  small  piece  of  some 
colored  cotton  fabric  (calico)  or  a  piece  of  litmus  paper. 

EXPERIMENT  39. — Combustion  of  chlorine  in  hydrogen. 
(Use  the  hood  in  this  experiment.) — Generate  chlorine 
from  coarsely  powdered  manganese  dioxide  and  strong 
hydrochloric  acid  in  a  wide  test-tube,  which  must  not  be 
more  than  one-third  filled  with  the  mixture.  The  tube  is 
clamped  upright  and  is  closed  by  a  perforated  cork  bear- 
ing a  straight  narrow  tube  about  20  cm.  long  ending  in  a 
jet.  Have  ready  a  bottle  of  hydrogen,  standing  in  water. 
Apply  a  gentle  heat  to  the  chlorine  generator,  and  when  the 
gas  over  the  liquid  is  green  and  chlorine  is  escaping  freely, 
light  the  hydrogen,  holding  the  bottle  inverted,  and  in- 
stantly slip  the  bottle  over  the  jet  of  chlorine.  Taken 
in  connection  with  the  burning  of  hydrogen  in  chlorine, 
what  does  the  result  show  ?  Test  the  product  in  the  bottle 
with  a  drop  of  ammonia  on  the  end  of  a  glass  rod. 

Fill  the  test-tube  with  cold  water  to  stop  the  produc- 
tion of  chlorine. 

Sum  up  all  the  evidence  you  have  obtained  thus  far 
bearing  upon  the  composition  of  hydrochloric  acid. 

EXPERIMENT  40. — Electrolysis  of  hydrochloric  acid. — 
Place  about  200  c.c.  of  strong  hydrochloric  acid  in  a 
beaker  and  dilute  it  with  about  J  of  its  volume  of  water. 
Add  a  handful  of  salt  and  stir  for  five  minutes  to  saturate 
the  liquid.  Pour  off  the  clear  liquid  and  introduce  it  into 


56  ELEMENTARY  CHEMISTRY 

the  apparatus  shown  in  Fig.  2,  Part  I,  which  you  have 
already  used  for  the  electrolysis  of  water.  The  salt  takes 
no  part  in  the  change.  Its  object  is  to  reduce  the  solubility 
of  the  chlorine  in  the  water  present. 

Allow  the  current  to  pass.  Allow  some  of  the  gases 
to  collect.  Notice  the  color,  and,  cautiously,  the  odor  of 
both.  Remember  that  a  colored  gas  may  appear  colorless 
in  a  thin  layer.  Look  along  the  tube  as  well  as  through  it. 
Try  to  burn  each  gas,  using  a  very  short  rubber  tube  with 
a  short  jet,  as  in  the  electrolysis  of  water.  Try  the  action 
of  both  gases  upon  paper  wet  with  a  solution  containing 
potassium  iodide  and  starch.  After  the  apparatus  has 
been  running  about  half  an  hour,  let  the  gases  collect 
for  a  time  and  measure  the  quantity  of  each.  Is  the  result 
exactly  what  you  expect  ?  If  not,  explain. 

EXPERIMENT  41. — Preparation  and  properties  of  hy- 
drochloric acid. — Prepare  hydrochloric  acid  from  commer- 
cial hydrochloric  acid  solution  and  strong  sulphuric  acid 
as  directed  on  p.  81,  Part  I.  A  flask  closed  by  a  doubly 
perforated  cork  carrying  a  separating  funnel  and  a  de- 
livery tube  is  sufficient.  The  sulphuric  acid  must  be 
allowed  to  drop  slowly  into  the  hydrochloric,  and  the 
delivery  tube  must  be  dry.  Collect  by  downward  dry 
displacement  in  bottles,  which  must  be  perfectly  dry. 
(Why?)  Keep  the  bottles  covered  during  the  collecting 
and  afterward. 

What  are  the  color  and  odor  of  the  gas  ?  Is  it  soluble 
in  water?  Bring  the  mouth  of  a  bottle  of  the  gas  near 
the  surface  of  water,  remove  the  glass  plate  and  then 
plunge  the  bottle  into  the  water.  Do  not  remove  the 
plate  under  the  water.  The  sudden  inrush  of  water  is 
likely  to  break  the  plate  and  cut  the  hand.  Wrap  the 
bulb  of  a  thermometer  in  damp  filter  paper  and  introduce 
it  into  a  bottle  of  hydrochloric  acid.  Is  there  any  altera- 
tion in  temperature?  Explain.  Does  the  gas  burn? 


CHLORINE  57 

Does  it  support  combustion  ?  Lower  a  lighted  candle  on  a 
wire  into  a  bottle  of  itv  In  another  bottle  place  a  strip  of 
red  and  a  strip  of  blue  litmus  paper.  (?)  Into  another 
bottle  throw  a  wad  of  filter  paper  wet  with  ammonia.  (?) 

EXPERIMENT  42. — Preparation  of  sodium  amalgam. — 
Weigh  out  roughly  on  the  trip-scales  in  a  dry  beaker  about 
200  grams  of  mercury.  Weigh  also  1.5-2  grams  of  so- 
dium and  cut  it  into  pieces  about  half  the  size  of  a  pea. 
Place  the  mercury  in  a  clean  dry  mortar,  add  one  piece  of 
sodium  and  press  it  under  the  mercury  with  the  pestle 
until  there  is  a  flash  and  a  little  smoke,  showing  that  the 
reaction  has  occurred.  Mercury  vapor  is  poisonous,  and 
it  is  well  therefore  to  do  this  under  the  hood.  Continue 
in  this  way  until  all  the  sodium  is  introduced.  The  amal- 
gam should  still  be  liquid.  If  it  is  solid,  too  much  sodium 
has  been  used  and  it  is  necessary  to  add  more  mercury  and 
mix  thoroughly  with  the  pestle.  Sodium  amalgam,  like 
sodium,  is  acted  upon  by  moist  air  and  must  be  preserved 
in  a  tightly  closed  bottle.  Its  chemical  action  is  simply 
that  of  the  sodium  it  contains,  but  the  presence  of  the 
mercury  makes  the  reactions  less  violent. 

EXPERIMENT  43. — Analysis  of  hydrochloric  acid. — 
Allow  10  c.c.  of  water  from  a  pipette  or  burette  to  run 
into  a  small  clean  test-tube  and  mark  the  level  of  the  water 
by  a  narrow  strip  of  gummed  label.  Dry  the  tube  thor- 
oughly with  paper  or  cloth,  not  with  the  flame.  Select 
a  straight  graduated  tube  of  about  50  c.c.  capacity  sealed 
at  one  end.  The  portion  nearest  the  open  end  will  not  be 
graduated.  Ascertain  its  capacity  by  taking  the  length 
from  the  end  of  the  graduation  to  the  end  of  the  tube  in 
compasses  or  with  a  strip  of  paper  and  transferring  it  to 
the  scale  on  the  tube.  Add  this  to  the  graduated  portion 
to  ascertain  the  total  capacity  of  the  tube. 

The  tube  must  be  perfectly  dry.  Fill  it  by  downward 
displacement  with  hydrochloric-acid  gas,  generated  as  in 


58  ELEMENTARY  CHEMISTRY 

Experiment  41,  and  dried  by  passing  through  a  U-shaped 
tube  containing  bits  of  broken  glass  tubing  wet  with  strong 
sulphuric  acid.  Since  the  presence  of  air  will  spoil  the 
result,  the  generator  must  be  run  until  all  air  is  expelled 
from  it  and  the  graduated  tube — five  minutes  or  more. 
The  use  of  the  hood  will  save  discomfort.  The  exit  tube 
must  run  to  the  bottom  of  the  graduated  tube  and  must 
be  withdrawn  gradually  while  the  generator  is  still  run- 
ning, otherwise  it  would  leave  a  vacancy  which  would  be 
filled  with  air.  Have  at  hand  10  c.c.  of  sodium  amalgam 
in  the  measuring  test-tube  and  at  once  introduce  it  and 
cover  the  tube  tightly  with  the  thumb.  Let  the  sodium 
amalgam  fall  the  length  of  the  tube  a  dozen  times  or 
more  to  make  the  action  complete,  bring  the  open  end  of 
the  tube  under  water  in  a  cylinder,  remove  the  thumb, 
lower  the  tube  until  the  level  of  the  water  is  the  same 
inside  and  out,  and  clamp  it.  Allow  to  stand  five  min- 
utes and  read  the  volume  of  the  gas. 

Cover  the  tube  with  the  thumb,  remove  it  from  the 
water,  invert  it  and  apply  a  flame  to  the  residual  gas.  (?) 
Calculate  thus: 

Volume  of  tube  to  end  of  graduated  portion 50  c.c. 

Capacity  of  ungraduated  portion 12  " 

Total  capacity 62  " 

Since  the  10  c.c.  of  sodium  amalgam  introduced 
expel  10  c.c.  of  gas,  the  volume  of  hydrochloric 
acid  taken  was 52  c.c. 

The  tube  at  the  end  of  the  experiment  contained  26.4  " 

Hence  the  proportion  of  hydrogen  by  volume  contained 

Q£»      A 

in  hydrochloric  acid  is  -  ^-,  or  about  J. 

O-c 

EXPERIMENT  44. — Action  of  sulphuric  acid  on  salt.— 
Place  a  handful  of  salt  in  a  wide-mouth  bottle  and  allow 
a  little  strong  sulphuric  acid  to  drop  upon  it.  Cover  the 


CHLORINE  59 

Uju  Tio 

bottle.     Is  heat  produced?     Sink  a  lighted  candle  in  the 

bottle.  (?)  Blow  across  the  mouth  of  it.  Hydrochloric 
acid  causes  fumes  in  moist  air,  because  it  causes  the  water- 
vapor  to  condense,  forming  a  solution  of  hydrochloric 
acid.  Hold  a  drop  of  ammonia  in  the  bottle.  (?) 

What  substance  would  remain  in  the  bottle  if  the 
action  of  the  sulphuric  acid  was  complete?  Write  the  • 
equation.  Place  about  25  c.c.  of  water  in  a  beaker.  Stand 
the  beaker  in  a  pan  of  cold  water  and  slowly  pour  in  50  c.c. 
strong  sulphuric  acid  stirring  constantly.  Let  the  liquid 
cool  completely  and  pour  it  slowly  upon  about  30  grams 
of  salt  in  a  small  flask  provided  with  a  well-fitting  cork 
and  delivery  tube.  This  tube  must  be  perfectly  dry,  for 
any  water  will  absorb  the  gas.  Apply  a  gentle  heat.  Be 
careful  not  to  let  the  liquid  boil  over.  Collect  the  hydro- 
chloric-acid gas  in  dry  test-tubes  and  bottles.  The  de- 
livery tube  must  reach  to  the  bottom  of  the  bottle.  Is  it 
soluble  in  water  ?  Try  it  by  placing  a  test-tube  filled  with 
it  mouth  downward  in  water.  How  does  it  affect  red  and 
blue  litmus  paper?  Does  it  burn?  Does  it  support  the 
combustion  of  a  candle?  Pour  some  ammonia-water  on 
a  wad  of  filter  paper  and  throw  it  into  a  bottle  of  the 
gas.  (?)  What  is  left  in  the  generating  flask? 

Pour  the  material  in  the  flask  down  the  sink,  running 
abundant  water  at  the  same  time,  and  again  put  in  salt 
and  dilute  sulphuric  acid  in  exactly  the  same  way.  Pass 
the  gas  into  a  small  inverted  funnel  whose  rim  dips  into 
water  in  a  dish  and  whose  stem  is  connected  by  a  rubber 
tube  with  the  delivery  tube.  Continue  until  the  mate- 
rials are  exhausted.  Use  the  hydrochloric  acid  prepared 
in  this  way  for  Experiment  45. 

EXPERIMENT  45. — Action  of  hydrochloric  acid  upon 
metals. — Try  the  action  of  hydrochloric  acid  upon  small 
quantities  of  zinc,  magnesium,  aluminium,  and  copper. 
Place  the  metal  in  a  test-tube,  cover  it  with  water,  and 


60  ELEMENTARY  CHEMISTRY 

gradually  add  hydrochloric  acid  solution.  If  gas  is  pro- 
duced test  it  with  a  flame.  (?)  What  is  the  other  product 
of  each  case?  Try  the  action  of  hydrochloric  acid  upon 
a  fragment  of  marble. 

QUESTIONS 

Answer  these  questions  as  briefly  as  possible.  If  you  know  the 
equation  for  the  change  you  desire  to  explain  that  alone  will 
suffice,  but  never  invent  equations  or  formulas.  Remember  that  a 
formula  is  a  faithful  description  of  the  composition  of  a  real  com- 
pound. Mere  arbitrary  collections  of  symbols  are  worse  than 
meaningless — they  are  absurd.  Never  write  an  equation  in  your 
notes  unless  you  are  sure,  first,  that  every  formula  it  contains  is 
that  of  a  real  substance  ;  second,  that  the  number  of  symbols  of 
each  kind  is  the  same  on  both  sides  of  the  sign  of  equality  ;  third, 
that  the  equation  describes  a  process  which  you  know  from 
laboratory  or  lecture-table  experiments  to  be  a  real  chemical 
change.  Unless  the  equation  satisfies  these  requirements,  discard 
it,  and  describe  the  change  briefly  in  words.  Some  of  the  ques- 
tions can  be  answered  in  a  number  of  different  ways,  all  of  which 
are"  correct. 

1.  How  could  you  convert  the  hydrogen  of  hydrochloric  acid 
into  water  ? 

Answer 

Zn  +  2HC1  =  ZnCU  +  H9. 

Ha  +  O  =  H,O. 

In  words,  the  answer  would  be  :  Zinc  will  liberate  the  hydro- 
gen, forming  zinc  chloride.  The  hydrogen  can  then  be  burned  in 
the  air  to  water. 

2.  How  would  you  make  NaaSO4  from  NaCl  ? 

3.  How  could  you  prepare  water,  taking  the  hydrogen  from 
sulphuric  acid  and  the  oxygen  from  copper  oxide  ? 

4.  Prepare  water,  taking  the  hydrogen  from  sulphuric  acid 
and  the  oxygen  from  potassium  chlorate. 

5.  Prepare  arsenic  chloride,  taking  the  arsenic  in  the  free  state 
and  the  chlorine  from  sodium  chloride. 

6.  How  could  you  convert  the  oxygen  of  potassium  chlorate 
into  ozone  ? 

7.  Obtain  hydrochloric  acid,  taking  the  hydrogen  from  water. 

8.  Convert  the  oxygen  of  water  into  copper  oxide. 


CHLORINE  61 


PROBLEMS 

32.  What  weight  of  chlorine  can  be  obtained  by  heating  12.5 
grams  of  manganese  dioxide  with  an  excess  of  hydrochloric  acid  ? 

33.  How  much   manganese    dioxide  is  needed   to  make   25 
grams  of  chlorine  from  hydrochloric  acid  ? 

34.  How  much  aqueous  hydrochloric  acid  containing  20  per 
cent  HC1  is  required  to  liberate  100  grams  of  chlorine  with  man- 
ganese dioxide  ? 

35.  20  c.c.  of  chlorine  were  mixed  with  16  c.c.  of  hydrogen 
and  the  mixture  exploded.     What  volumes  of  what  gases  remained 
in  the  vessel  ? 

36.  Hydrogen  was  burned  in  chlorine  and  the  hydrochloric 
acid  produced  collected.     It  weighed  146  grams.     What  weights 
of  both  gases  had  been  consumed  ? 

37.  50  grams  of  sodium  hydroxide  are  dissolved  in  water  and 
the  solution  mixed  with  a  solution  containing  50  grams  of  pure 
hydrochloric  acid.     What  substances,  and  how  much,  will  be  con- 
tained in  the  resulting  liquid  ?     Do  not  calculate  water. 

38.  32.75  grams  of  zinc  are  dissolved  in  hydrochloric  acid. 
What  weights  of  zinc  chloride  and  hydrogen  are  produced  ? 

39.  40  grams  of  magnesium  are  allowed  to  dissolve  in  hydro- 
chloric acid.     Calculate  the  weights  of  hydrogen  and  magnesium 
chloride  produced. 


CHAPTER   XI 

CHLORIDES-COMPOUNDS  OF  CHLORINE  CONTAINING 
OXYGEN 

EXPERIMENT  46. — Preparation  of  chlorides. — Place 
about  1  gram  of  zinc  oxide  in  a  test-tube,  cover  with 
water,  and  add  hydrochloric  acid.  If  necessary,  heat  gen- 
tly, but  not  to  boiling.  What  is  produced  ?  Is  gas  evolved  ? 
What  becomes  of  the  hydrogen  of  the  acid  ?  Repeat,  using 
magnesium  oxide.  Write  the  equations. 

Make  a  strong  solution  of  sodium  hydroxide  and  slow- 


62  ELEMENTARY  CHEMISTRY 

ly  add  an  excess  of  hydrochloric  acid,  stirring  constantly. 
Write  the  equation. 

Examine  a  fragment  of  marble.  Coarsely  powder  a 
little.  Is  it  soluble  in  water  ?  Add  some  hydrochloric  acid 
to  the  water.  (?)  When  the  action  is  over,  evaporate  the 
liquid  to  dry  ness  in  a  dish.  Is  the  residue  soluble  in 
water?  Is  it  marble?  Write  the  equation. 

State  four  methods  by  which  the  salt  of  a  given  metal 
and  a  given  acid  can  be  prepared,  and  give  an  example  of 
each. 

EXPERIMENT  47. — Insoluble  and  slightly  soluble  chlo- 
rides.— To  a  dilute  solution  of  silver  nitrate,  AgN03,  add 
a  little  hydrochloric  acid.  The  precipitate  is  silver  chlo- 
ride, AgCl.  Write  the  equation.  Divide  the  liquid  con- 
taining the  precipitate  into  three  parts.  Expose  one  por- 
tion to  sunlight,  or  the  brightest  light  attainable.  With 
the  second  ascertain  whether  silver  chloride  is  soluble  in 
nitric  acid.  With  the  third  investigate  its  solubility  in 
ammonia. 

Mix  a  few  drops  of  silver  nitrate  solution  with  some 
solution  of  sodium  chloride.  What  is  the  precipitate? 
Prove  it.  Any  soluble  chloride  would  produce  the  same 
effect  upon  silver  nitrate. 

Precipitate  a  little  solution  of  mercurous  nitrate, 
Hg2(N03)2,  with  hydrochloric  acid.  The  precipitate  is  ^ 
mercurous  chloride,  Hg2Cl2.  Divide  the  liquid  containing 
the  precipitate  into  two  parts.  With  one  ascertain  the 
action  of  strong  light.  Expose  for  some  time.  Treat  the 
other  with  ammonia.  What  are  the  results?  How  can 
mercurous  chloride  be  distinguished  from  silver  chloride? 

Make  a  solution  of  lead  nitrate,  Pb(N03)2,  and  add 
hydrochloric  acid  to  it.  The  precipitate  is  lead  chloride, 
PbCl2.  Write  the  equation.  Heat  the  liquid  containing 
the  lead  chloride  to  boiling  and  stand  it  aside  to  cool.  Is 
lead  chloride  soluble  in  cold  water  ?  In  hot  water  ?  How 

U 


CHLORIDES  63 

could  lead  chloride  be  separated  from  silver  chloride? 
From  mercurous  chloride  ?  Devise  a  method  of  separating 
a  mixture  of  the  three  chlorides. 

EXPERIMENT  48. — Chlorine  peroxide. — Make  the  ex- 
periments described  in  paragraph  112,  Part  I.  Use  a 
cylinder  of  not  more  than  200  c.c.  capacity  and  not  more 
than  0.5  gram  of  finely  powdered  potassium  chlorate. 
Cover  the  bottom  of  the  cylinder  with  sulphuric  acid  and 
introduce  the  chlorate  gradually. 

EXPERIMENT  49. — Potassium  chlorate.  Potassium  hy- 
droxide must  not  be  touched  with  the  hands.  Use  paper 
or  forceps  in  handling  it.  Make  this  experiment  under 
the  hood. — Dissolve  about  40  grams  of  potassium  hydrox- 
ide in  about  100  c.c.  of  water  in  a  dish  and  pass  chlorine 
into  the  liquid  through  an  inverted  funnel  whose  rim  dips 
under  the  liquid.  From  time  to  time  test  the  liquid  -with 
a  strip  of  red  litmus  paper,  and  when  the  latter  is  no 
longer  turned  blue  (?)  stop  the  chlorine.  Pour  off  the 
liquid  from  the  white  crystals  in  the  dish  and  dissolve  the 
latter  in  the  smallest  possible  quantity  of  hot  water.  If 
anything  remains  undissolved,  remove  it  by  filtering  the 
hot  liquid,  then  let  the  liquid  cool.  Collect  the  product 
which  separates  on  a  filter,  and  dry  it  between  layers  of 
filter  paper.  Examine  it  and  compare  its  appearance  with 
that  of  potassium  chlorate.  What  is  the  effect  of  heat  upon 
it?  Use  a  tube  sealed  at  one  end  and  test  the  gas  given 
off  by  the  spark.  Unless  your  substance  was  perfectly  dry, 
water  will  appear  in  the  sealed  tube,  but  it  is  not  a  product 
of  the  chemical  change  and  can  be  disregarded. 

What  is  the  action  of  sulphuric  acid  on  the  substance 
you  have  made  ?  Use  very  small  quantities. 

How  is  potassium  chlorate  made  at  present,  and  what 
are  its  chief  uses? 


64  ELEMENTARY   CHEMISTRY 

QUESTIONS 

1.  Why  is  no  gas  given  off  when  an  oxide  or  a  hydroxide  is 
treated  with  hydrochloric  acid  ? 

2.  What  is  the  real  meaning  of  the  expression  '  *  insoluble  "  ? 

3.  How  could   sodium  chloride    be  converted  into    sodium 
nitrate  ? 

4.  Does  chlorine  peroxide,  ClOa,  contain  more  or  less  energy 
than  chlorine  and  oxygen  separately  ?     How  do  you  know  ? 

5.  Construct,  in  the  form  of  a  table,  a  brief  comparison  of  the 
four  elements  already  studied. 

PROBLEMS 

40.  Calculate  the  percentage  composition  of  bleaching  powder, 
CaOCl,. 

41.  Determine  the  name  and  formula  of  a  compound  having 
the  following  composition : 

Sodium    21.60  per  cent. 

Chlorine 33.33        " 

Oxygen 45.07        " 

42.  25  grams  of  pure  marble  are  dissolved  in  hydrochloric  acid. 
What  are  the  products  and  how  much  of  each  is  produced.     How 
much  HC1  is  consumed  ? 


CHAPTER   XII 
No  experiments. 

QUESTIONS 

1.  What  are  the  two  possibilities  regarding  the  nature  of 
matter  ? 

2.  What  is  a  molecule  ?    Is  a  molecule  a  thing  or  an  idea  ? 
Why? 

3.  What  is  the  evidence  which  causes  us  to  consider  the  mole- 
cule to  be  composed  of  smaller  particles  ? 

4.  Explain  precisely  what  is  meant  by  the  term  atomic  weight. 

5.  State  the  chemical  laws  in  the  language  of  the  atomic  theory. 

6.  Suppose  that  some  one  should  succeed  in  transforming  lead 


THE  ATMOSPHERE— NITROGEN  65 

into  gold  -  a  highly  improbable  supposition  :  what  changes  in  the 
atomic  theory  would  have  to  be  made  to  explain  the  discovery  ? 


CHAPTER     XIII 

THE  ATMOSPHERE-NITROGEN 

EXPERIMENT  49a. — Analysis  of  the  air  by  means  of 
phosphorus. — Phosphorus  catches  fire  spontaneously,  and 
to  leave  it  lying  around  or  put  away  any  of  it  under  the 
desk  would  be  likely  to  cause  the  destruction  of  the  build- 
ing. Burns  made  with  it  are  poisoned  wounds  which  heal 
with  great  difficulty.  It  must  be  kept,  cut,  and  handled 
under  water,  and  the  hands  must  be  freed  from  it  before 
they  are  brought  into  the  air.  Be  careful  not  to  get  it 
behind  the  nails.  If  it  should  catch  fire  while  you  are 
working  with  it  pour  water  upon  it. 

The  method  is  to  absorb  the  oxygen  from  a  measured 
volume  of  air,  confined  over  water,  by  means  of  phos- 
phorus. The  residual  gas,  consisting  chiefly  of  nitrogen, 
is  again  measured.  The  experiment  is  shown  in  Fig.  27, 
Part  I.  The  water  used  must  have  stood  in  the  laboratory 
a  day  or  two  so  as  to  acquire  the  temperature  of  the  room. 
The  jar  should  be  wide,  so  as  to  allow  freedom  to  the 
hand. 

Fill  the  graduated  tube  with  water,  close  the  end  with 
the  thumb  and  put  it,  open  end  down,  into  the  jar  of  water. 
Catch  hold  of  the  upper  part  of  the  tube  with  a  paper 
holder — not  with  the  hand — and  allow  20  c.c.  to  30  c.c.  of 
air  to  enter  the  tube.  Hold  the  tube  so  that  the  level  of 
the  water  inside  and  outside  is  the  same,  and  measure  the 
volume  of  air.  Repeat  the  measurement  to  be  sure  there 
is  no  error. 

Now  bend  a  wire  to  the  shape  shown  in  the  cut  and 


66  ELEMENTARY  CHEMISTRY 

fasten  a  small  piece  of  phosphorus  to  it,  working  under 
water  in  a  tin  pan  or  other  large  vessel.  Transfer  the 
wire  to  the  jar  and  push  the  phosphorus  up  into  the  air 
in  the  graduated  tube,  taking  care  to  keep  the  open  end 
of  the  tube  constantly  under  water.  Let  the  apparatus 
stand  for  an  hour  and  examine  it.  If  the  water  has  reached 
the  phosphorus  raise  the  latter  so  as  to  keep  it  in  the  air. 
Next  morning  pull  the  phosphorus  down  into  the  water 
and  make  the  level  of  the  water  inside  and  outside  the 
same.  This  must  be  done  without  touching  the  -tube  with 
the  fingers.  Read  the  volume  of  the  residual  gas.  The 
loss  in  volume  is  oxygen.  Calculate  the  percentage  of 
oxygen  by  volume  in  the  air  thus : 

Volume  of  oxygen  X  100 

— ^ — = —      -  =  per  cent  of  oxygen. 
Volume  of  air 

Plunge  a  burning  splinter  into  the  gas  remaining  in 
the  tube. 

In  this  method  errors  result  from  changes  in  atmos- 
pheric pressure  and  temperature  between  the  two  measure- 
ments. If  a  more  accurate  result  is  desired,  a  thermometer 
must  be  placed  in  the  jar  ten  minutes  before  the  volume 
of  the  air  is  measured  in  the  first  place.  Temperature  and 
barometer  pressure  must  be  read  and  recorded  and  the 
volume  of  the  air  reduced  to  standard  conditions  by  the 
method  explained  in  the  Appendix.  Temperature  and 
pressure  must  again  be  read  at  the  time  the  residual  gas 
is  measured  and  its  volume  also  reduced.  Then  this  re- 
duced volume  is  subtracted  from  the  other  and  the  cal- 
culation made  as  before. 

EXPERIMENT  49Z>. — Analysis  of  the  air.  Alternative 
method.1 — Use  a  graduated  tube  holding  50  c.c.,  like  that 

1  This  method  is  at  least  as  accurate  as  that  described  in  49a,  and 
much  more  rapid.  It  has  the  advantage  of  not  requiring  the  use  of 
phosphorus. 


THE  ATMOSPHERE— NITROGEN  67 

employed  in  Experiment  43.  Ascertain  the  capacity  of 
the  ungraduated  portion.  Have  ready  a  cork  which  fits 
the  tube  tightly.  If  a  wooden  cork,  it  must  be  well  rolled 
under  the  foot.  A  rubber  cork  without  any  perforation 
is  best.  Make  a  measuring  test-tube  to  hold  10  c.c.,  as 
described  in  Experiment  43.  It  need  not  be  dry. 

Dissolve  a  small  piece  of  potassium  hydroxide  in  about 
15  c.c.  of  water  in  a  beaker  and  add  to  the  liquid  about 
0.5  gram  of  pyrogallic  acid.  Use  the  solution  at  once.  It 
spoils  rapidly.  Take  10  c.c.  of  this  liquid  in  your  measur- 
ing test-tube,  pour  it  into  the  graduated  tube,  and  in- 
stantly cork  the  tube  tightly.  Make  the  liquid  flow  from 
one  end  of  the  tube  to  the  other  twenty  times  or  more. 
Place  it  in  a  deep  cylinder  of  water,  cork  down,  remove 
the  cork,  allow  to  stand  five  minutes,  equalize  the  levels 
without  touching  the  part  of  the  tube  containing  the  nitro- 
gen with  the  hand  and  read  the  volume.  Calculate  thus : 

Capacity  of  tube  to  end  of  graduations 50  c.c. 

Capacity  of  ungraduated  portion 9  " 

Total  capacity 59  " 

Since  10  c.c.  of  liquid  was  introduced,  the  volume  of 
air  taken  was  59  —  10,  or  49  c.c. 

Volume  of  gas  left  in  the  tube,  38.5  c.c. 
The  loss  in  volume  is  oxygen.     This  is 

49  -  38.5  =  10.5  c.c., 
and  the  percentage  of  oxygen  is 

10.5  X  100 

=  21.4  per  cent. 

EXPERIMENT  50. — Eudiometric  analysis  of  air. — In 
this  experiment  the  U-shaped  eudiometer  described  in  Ex- 
periment 9  is  used.  If  no  Kipp  generator  is  available, 
have  ready  a  wide  test-tube  with  a  delivery  tube  ending  in 
a  short  piece  of  rubber  tubing.  This  test-tube  should  con- 
27 


68  ELEMENTARY  CHEMISTRY 

tain  a  little  zinc  covered  with  water  and  is  to  be  used  as 
a  source  of  hydrogen. 

Clamp  the  eudiometer  (Fig.  9)  vertically  and  fill  both 
limbs  of  it  with  mercury.  Open  the  stop-cock  S,  place  a 
dry,  clean  beaker  under  C  and  allow  mercury  to  run  out 
until  about  20  c.c.  of  air  has  entered  the  tube.  Close  C 
and  measure  and  record  the  volume  of  the  air.  Verify 
the  measurement.  Turn  8  so  that  the  air  in  the  eudiom- 
eter is  isolated  and  so  that  gas  passed  in  at  T  will  escape 
at  E.  Pass  hydrogen  through  the  stop-cock  for  three 
minutes,  allowing  it  to  escape  into  the  air.  Then  pass 
not  more  than  15  c.c.  nor  less  than  10  c.c.  of  hydrogen 
into  the  eudiometer,  turn  the  stop-cock  so  as  to  close  the 
eudiometer,  and  at  once  remove  the  generator.  Equalize 
the  mercury  levels  in  both  limbs  by  running  out  mercury 
carefully  from  C.  If  too  much  mercury  is  run  out,  pour 
some  in  at  the  top  of  the  open  limb  and  again  equalize. 
Read  off  the  volume  and  record  it  as  air  -f-  hydrogen. 

Press  the  thumb  tightly  on  the  open  end  of  the  eudi- 
ometer and  pass  the  spark  from  an  induction-coil  excited 
by  three  Edison-Lalande  cells,  two  bichromate  cells,  or 
in  any  other  suitable  way.  What  is  the  cause  of  the  ex- 
plosion? Pour  mercury  in  the  open  limb  until  the  level 
is  slightly  higher  than  that  in  the  other,  and  let  the  ap- 
paratus stand  for  five  minutes.  (?)  Equalize  the  levels 
exactly  by  running  out  mercury  from  (7,  and  read  and 
record  the  "  volume  after  explosion."  Calculate  as  in  the 
following  example: 

Volume  of  air  taken 25     c.c. 

Volume  of  air  -f-  hydrogen 39       " 

Volume  after  explosion 23.4    " 

The  loss  in  volume  after  explosion  must  be  15.6    " 

Now  this  contraction  is  due  to  the  disappearance  of 
oxygen  and  hydrogen  to  form  water,  and  in  this  reaction 


THE  ATMOSPHERE— NITROGEN  69 

one  volume  of  oxygen  and  two  of  hydrogen  disappear. 
Hence  the  volume  of  the  oxygen  in  the  air  taken  must  be 

-  =  5.2  c.c.,  and  the  percentage  must  be  — — ^ = 

o  &o 

20.8  per  cent. 

This  method  of  analyzing  the  air  is  employed  in  actual 
work,  since  it  is  rapid,  and,  with  proper  care,  accurate. 

EXPERIMENT  51. — Absorption  of  the  oxygen  of  the  air 
by  metals. — The  method  is  to  pass  air  over  a  column  of 
copper  filings  heated  nearly  to  redness  in  a  tube  similar 
to  that  used  in  Experiment  10,  Fig.  11.  The  tube  should 
be  about  30  cm.  (12  inches)  long,  and  about  20  cm.  (8 
inches)  of  it  should  be  heated  nearly  to  redness  by  a  wing- 
top  burner  with  several  chimneys.  The  copper  filings  are 
held  in  place  by  loose  plugs  of  asbestos,  and  after  being 
filled  the  tube  should  be  held  horizontal  and  tapped  gently 
on  the  table  to  make  a  channel  along  the  upper  portion. 


PIG. 

One  end  of  the  tube  is  connected  with  a  bottle  out  of  which 
air  is  expelled  by  water,  which  is  allowed  to  drop  into  the 
bottle  from  a  separating  funnel  which  passes  through  the 
stopper.  The  air  passes  over  the  hot  copper,  which  absorbs 
the  oxygen,  and  the  nitrogen  is  collected  over  water.  The 
apparatus  is  shown  in  Fig.  21.  The  bottle  from  which 
the  air  is  expelled  and  that  in  which  the  nitrogen  is  col- 


70  ELEMENTARY  CHEMISTRY 

lected  should  be  of  the  game  size  and  shape.  The  speed 
of  the  current  of  air  depends  upon  the  rate  at  which  the 
water  is  allowed  to  fall  upon  the  separating  funnel.  It 
should  be  as  slow  as  time  permits. 

What  change  takes  place  in  the  copper  ?  At  which  end 
of  the  tube  does  the  change  begin?  About  what  fraction 
of  the  air  disappears? 

When  the  experiment  is  finished,  slip  a  glass  plate  over 
the  bottle  of  nitrogen  and  remove  it  from  the  water. 
Plunge  a  burning  piece  of  wood  or  a  lighted  candle  into 
the  gas.  ( ?)  What  is  the  effect  of  the  nitrogen  of  the 
air  upon  combustion  processes,  and  how  is  this  effect 
caused  ? 

Eepeat  the  whole  experiment,  using  iron  filings  instead 
of  copper.  Explain  exactly  what  takes  place. 

EXPERIMENT  52. — Expose  some  clear  lime-water  in  a 
beaker  to  the  air  for  some  time.  Look  for  a  change  at  the 
surface  of  the  lime-water.  What  does  the  result  prove? 
What  are  the  sources  of  the  carbon  dioxide  of  the  air  and 
what  becomes  of  it? 

EXPERIMENT  53. — Nitrogen. — Put  about  5  grams  of 
ammonium  chloride  and  the  same  quantity  of  powdered 
sodium  nitrite,  NaN02,  in  a  small  flask  and  add  enough 
water  to  make  a  thin  paste.  Set  up  the  apparatus  as  shown 
in  Fig.  28,  Part  I.  Apply  a  gentle  heat  and  collect  the 
nitrogen  over  water.  Stop  heating  as  soon  as  the  reaction 
begins  or  it  will  become  too  violent.  If  the  evolution  of 
gas  threatens  to  become  too  energetic,  immerse  the  gen- 
erating flask  for  an  instant  only  in  a  pan  of  cold  water, 
which  should  be  in  readiness.  This  will  immediately  quiet 
it.  Long  immersion  will  cause  water  to  flow  back  into 
the  flask  and  spoil  the  experiment. 

What  are  the  physical  properties  of  the  substance  you 
have  prepared  ?  Does  it  burn  ?  Does  it  support  the  com- 
bustion of  a  candle  ? 


THE  ATMOSPHERE— NITROGEN  71 


QUESTIONS 

1.  What  constituents  of  the  air  are  essential  to  life,  and  why  ? 

2.  What  are  the  reasons  for  the  belief  that  the  air  is  a  mixture, 
not  a  compound  ? 

3.  Describe  an  experiment  which  proves  that  the  air  contains 
the  vapor  of  water. 

4.  Nitrogen  is  lighter  than  oxygen  and  water- vapor  only  about 
half  as  dense.     On  the  other  hand,  carbon  dioxide  is  much  denser 
than  oxygen.     Why  is  it  that  all  these  substances  remain  uniformly 
mixed  in  the  air  and  do  not  separate   into  layers  according  to 
density  ? 

5.  What  is  the  effect  of  the  nitrogen  of  the  air  upon  combus- 
tion processes,  and  why  ? 

6.  Is  there  any  difference  between  nitrogen  obtained  from  the 
atmosphere   according  to   Experiment  51  and  that  prepared  by 
chemical  methods  as  in  Experiment  53  ?     Why  ? 

7.  What  elements  are  most  abundant  in  the  bodies  of  animals 
and  plants  ? 

PKOBLEMS 

43.  16.7  c.c.  of  air  were  confined  over  mercury  in  a  eudiome- 
ter and   enough   hydrogen   added  to   make  the  volume  30  c.c. 
After  explosion,  the  volume  was  19.5  c.c.     What  percentage  of 
oxygen  by  volume  did  the  air  contain  ? 

44.  20  c.c.  of  air  are  mixed  with  10  c.c.  of  hydrogen  and  the 
spark  is  passed.     After  the  explosion,  what  volumes  of  what  gases 
remain  in  the  tube  ? 

In  solving  this  problem  assume  that  the  air  contains  21  per 
cent  by  volume  of  oxygen. 

45.  2  liters  of  air  were  passed  over  hot  copper.     The  increase 
in  weight  of  the  copper  was  .6  gram.     What  was  the  percentage 
of  oxygen  by  weight  in  the  air  ? 

Assume  the  weight  of  1  liter  of  air  to  be  1.293  grams. 

46.  What  weight  of  nitrogen  can  be  obtained  by  heating  13.8 
grams  of  sodium  nitrite,    NaNO2,  with  the  required  quantity  of 
ammonium  chloride,  NH4C1  ?     How  much  ammonium  chloride  is 
needed  ?     What  weights  of  salt  and  of  water  are  formed  ? 


72  ELEMENTARY  CHEMISTRY 

CHAPTER   XIV 

AMMONIA 

EX^RIMENT  54. — Preparation  of  ammonia  gas. — Place 
about  ;fe0  c.c.  of  ammonium  hydroxide  in  a  ^Qfr  c.c.  flask. 
Arrange  the  apparatus  as  shown  in  Fig.  30,  Part  I.  A 
little  mercury  should  be  placed  in  the  bend  of  the  safety- 
tube.  A  piece  of  asbestos  board  must  be  placed  under  the 
flask.  Use  glass  tubing  as  far  as  possible  in  constructing 
the  apparatus.  A  rubber  cork  can  be  used  in  the  flask. 

Apply  a  gentle  heat  and  collect  the  ammonia  by  up- 
ward displacement  in  bottles,  which  must  be  absolutely 
dry..  Why?  The  tube  must  run  up  to  the  bottom  of  the 
inverted  bottle.  From  time  to  time  hold  a  drop  of  hy- 
drochloric acid  on  a  glass  rod  to  the  mouth  of  the  bottle. 
Dense  white  fumes  indicate  that  the  latter  is  full  and  am- 
monia is  escaping.  Collect  three  bottles  of  the  gas  and 
stand  them  on  glass  plates  in  inverted  position.  Hold  the 
exit  tube  of  your  apparatus  to  the  burner  flame.  How 
does  the  gas  behave  as  regards  combustion?  Does  it  pro- 
duce a  continuous  flame?  Can  it  be  made  to  do  so? 
What  are  the  products  of  its  combustion? 

Plunge  the  mouth  of  a  bottle  filled  with  the  gas  under 
water  and  move  it  so  as  to  bring  the  water  in  contact  with 
the  gas.  (?)  Remove  the  glass  plate  before  placing  the 
mouth  of  the  bottle  under  the  water. 

Slowly  introduce  a  lighted  candle  held  on  a  wire  into 
an  inverted  bottle  of  the  gas,  as  you  did  with  hydrogen. 
Does  it  support  combustion?  How  does  it  behave  on  the 
instant  of  contact  with  the  candle-flame  ? 

Fill  a  dry  bottle  with  hydrochloric-acid  gas  by  down- 
ward displacement.  The  gas  can  be  made  by  gently  heat- 
ing a  little  strong  hydrochloric  acid  placed  in  a  test-tube 
with  a  cork  bearing  a  delivery  tube  bent  downward.  Bring 


AMMONIA  73 

the  bottle  mouth  to  mouth  with  a  bottle  of  ammonia  gas. 
Let  the  bottles  stand  until  the  product  has  settled  on  the 
sides.  Scrape  some  of  it  out  and  heat  it  carefully  in  a 
porcelain  crucible  or  on  a  piece  of  platinum  foil.  (?) 
What  is  it? 

EXPERIMENT  55. — Ammonium  salts. — Dilute  some  am- 
monium hydroxide  (ammonia  water)  with  about  three 
times  its  volume  of  water  in  a  beaker.  How  does  the 
liquid  affect  red  and  blue  litmus  paper  ?  Does  it  act  upon 
the  litmus  in  the  same  way  as  caustic  soda? 

Dilute  about  5  c.c.  of  hydrochloric  acid  with  three 
times  its  volume  of  water  in  a  dish  and  investigate  its 
behavior  with  both  kinds  of  litmus  paper.  All  acids  affect 
litmus  paper  in  the  same  way  as  hydrochloric  acid.  Am- 
monium hydroxide  and  sodium  hydroxide  are  bases.  All 
soluble  bases  affect  litmus  paper  in  the  same  way. 

Now  add  your  dilute  ammonium  hydroxide,  drop  by 
drop,  to  the  hydrochloric  acid  in  the  dish,  stirring  con- 
stantly. From  time  to  time  dip  small  pieces  of  blue  lit- 
mus paper  into  the  liquid.  When  the  latter  is  no  longer 
turned  red  you  will  find  that  a  piece  of  red  litmus  paper 
is  turned  blue.  At  the  same  time  the  liquid  acquires  a 
faint  odor  of  ammonia.  Why? 

Evaporate  the  liquid  in  the  dish  to  dryness.  Examine 
the  residue.  Is  it  ammonium  chloride?  Heat  some  of 
it  on  a  piece  of  platinum  foil  or  in  a  crucible.  Write  the 
equation. 

Repeat  the  whole  experiment,  using  nitric  acid  instead 
of  hydrochloric.  Write  the  equation. 

Dissolve  some  ammonium  chloride  in  water  in  a  test- 
tube.  Add  a  fragment  of  sodium  hydroxide  to  the  liquid 
and  heat  gently.  Notice  the  odor  of  the  gas  given  off. 
Hold  a  fragment  of  red  litmus  paper  in  the  gas.  Repeat, 
using  potassium  hydroxide  instead  of  sodium  hydroxide. 
In  the  same  way  investigate  the  behavior  of  ammonium 


74  ELEMENTARY   CHEMISTRY 

nitrate  with  sodium  hydroxide  and  potassium  hydroxide. 
Notice  that  strong  bases  liberate  ammonia  from  ammo- 
nium salts.  Write  equations  for  the  four  chemical  changes 
which  have  taken  place. 

Study  the  effect  of  heat  upon  small  quantities  of  am- 
monium nitrate,  chloride,  and  sulphate.  Use  a  porcelain 
crucible  or  a  piece  of  platinum  foil. 

EXPERIMENT  56. — Dilute  about  30  c.c.  of  ammonia 
with  about  ten  times  its  volume  of  water,  in  a  flask.  Shake 
up  the  liquid  with  salt  until  no  more  dissolves.  Fill  the 
apparatus  employed  in  the  electrolysis  of  water  (Fig.  6) 
with  the  clear  liquid  and  pass  the  current.  What  two 
gases  collect?  Try  to  burn  each.  Close  the  stop-cocks 
for  a  time  and  ascertain  the  relative  quantities  of  the 
gases.  If  the  apparatus  is  not  graduated,  measure  the 
length  of  the  columns  of  gas  with  a  meter  scale. 

In  this  experiment  the  salt  is  added  only  to  assist  in 
conducting  the  current  through  the  liquid,  and  thus  make 
the  process  more  rapid. 

QUESTIONS 

1.  What  does  the  presence  of  ammonia  indicate  with  regard  to 
the  fitness  of  a  sample  of  water  for  drinking  ? 

2.  Explain  the  fact  that  the  air  of  stables  often  smells  strongly 
of  ammonia. 

3.  Remembering  that  the   density   of  nitrogen  referred  to 
hydrogen  is  14,  and  that  the  symbol  N  means  14  parts  of  nitrogen 
by  weight,  deduce  from  Experiment  56  the  formula  of  ammonia. 

4.  Devise  two  methods  of  distinguishing  an  ammonium  salt 
from  a  sodium  or  a  potassium  salt. 

5.  How  would  you  convert  the  hydrogen  of  ammonia  into 
water  ? 

6.  How  would  you  convert  the  hydrogen  of  water  into  am- 
monia ?     (See  paragraph  139,  Part  I.) 

7.  How   would   you   convert   free   nitrogen   into  ammonium 
nitrate  ?    (See  p.  105,  Part  I.) 


AMMONIA  75 

8.  What  are  some  of  the  uses  of  ammonia  water  ?  Liquid 
ammonia  is  largely  employed  for  one  purpose  only.  What  is  it  ? 

0.  How  could  you  make  ammonia,  taking  the  hydrogen  from 
sulphuric  acid  and  the  nitrogen  from  ammonium  nitrite,  NH4NOa  ? 

10.  Explain  exactly  what  is  meant  by  a  radical. 

PKOBLEMS 

47.  856  grams  of  ammonium  chloride  are  heated  with  sodium 
hydroxide.     How  much  ammonia  by  weight  escapes  ? 

48.  Calculate  the  percentage  composition  of 

a.  Ammonium  chloride,  NH4C1. 
5.  Ammonium  nitrate,  NH4NC>3. 

c.  Ammonium  sulphate,  (NH4)2SO4. 

d.  Ammonium  hydroxide,  NH4OH. 

49.  When  a  stream  of  electric  sparks  is  passed  through  am- 
monia it  is  decomposed,  two  volumes  yielding  one  volume  of  nitro- 
gen and  three  volumes  of  hydrogen.     What  volumes  of  nitrogen 
and  hydrogen  are  formed  when  300  c.c.  of  ammonia  are  treated  in 
this  way  ? 

50.  100  c.c.    of  ammonia  are   decomposed   by  a  stream  of 
sparks,     (a)  What  volume  of  oxygen  must  be  added  to  the  result- 
ing mixture  to  combine  with  the  hydrogen  and  produce  water  ? 
(5)  After  the  water  has  condensed,  what  gas  will  remain  in  the 
tube,  and  how  much  ? 

51.  32  c.c.  of  ammonia  are  decomposed  by  sparks,  50  c.c.  of 
oxygen  are  added,  and  the  mixture  is  caused  to  explode.     What 
volumes  of  what  gases  are  left  ? 

52.  An   unknown   volume   of   ammonia  is  decomposed  in  a 
eudiometer,  an  unknown  volume  of  oxygen  is  mixed  with  it,  and 
the   mixture  exploded.     After  the  explosion  the  contraction  in 
volume  is  18  c.c.  and  the  tube  still  contains  some  oxygen,     (a) 
What  volume  of  ammonia  was  taken  in  the  first  place,  and  (b) 
what  volume  of  nitrogen  was  left  in  the  tube  ? 

53.  Ammonium   chloride   is   heated   in  a  flask  with   sodium 
hydroxide  and  the  ammonia  passed  into  31.5  grams  of  pure  nitric 
acid.     How  much  ammonium  chloride  must  be  used  in  order  to 
convert  all  the  nitric  acid  into  ammonium  nitrate  ? 


76  ELEMENTARY  CHEMISTRY 

CHAPTER    XV 

COMPOUNDS  OF  NITROGEN  AND  OXYGEN 

EXPERIMENT  57. — Nitrous  oxide. — Fill  a  wide  test- 
tube  one-third  with  ammonium  nitrate.  Clamp  it  in  an 
inclined  position,  and  insert  a  perforated  cork  with  a 
delivery  tube.  It  is  advisable  to  pass  the  gas  through  an 
empty  wide  test-tube  closed  by  a  doubly  perforated  cork 
before  collecting  it.  Apply  a  gentle  heat  and  collect  the 
gas  over  warm  water.  The  evolution  of  gas  must  be  slow 
— a  bubble  or  two  a  second.  If  it  becomes  too  rapid  ex- 
plosions result.  This  can  be  easily  controlled  by  lowering 
or  removing  the  flame.  Stop  heating  and  remove  the 
cork  before  the  ammonium  nitrate  is  exhausted.  Explo- 
sions sometimes  occur  when  the  quantity  of  substance  be- 
comes small.  If  the  ammonium  nitrate  shows  signs  of 
giving  out  before  you  have  enough  gas,  disconnect,  add 
more  of  it  directly  to  the  liquid  in  the  test-tube,  and 
resume  heating  cautiously.  Is  there  evidence  of  the  for- 
mation of  any  product  besides  the  gas?  What? 

Collect  three  bottles  and  two  test-tubes  full  of  the  gas. 
Use  one  bottle  to  determine  color,  odor,  and  taste.  (?) 
Try  the  spark  test  in  a  test-tube.  Ascertain  if  the  gas  is 
soluble  in  water.  Plunge  a  lighted  candle  into  a  bottle  of 
it.  (?)  Set  fire  to  some  sulphur  in  a  deflagrating  spoon, 
and  at  the  instant  the  sulphur  begins  to  burn  plunge  it 
into  the  third  bottle  of  nitrous  oxide.  It  should  be  extin- 
guished. Cover  the  bottle,  and  heat  the  sulphur  until  it 
burns  vigorously.  Plunge  it  again  into  the  gas.  (?)  Is 
nitrous  oxide  easy  or  difficult  to  decompose  into  its  ele- 
ments? Which  has  probably  the  higher  temperature,  the 
candle-flame  or  the  flame  of  burning  sulphur? 

EXPERIMENT  58.— Nitric  oxide.— Make  no  attempt  to 
ascertain  the  odor  of  nitric  oxide.  It  must  not  be  inhaled, 


COMPOUNDS  OF  NITROGEN  AND  OXYGEN          77 

and  experiments  with  it  should  be  carried  out  under  the 
hood,  if  possible.  If  you  do  not  work  under  the  hood, 
keep  the  bottles  of  the  gas  covered  during  combustions  and 
let  the  products  escape  out  of  the  windows. 

Nitric  oxide  is  made  in  the  same  apparatus  which  is 
used  for  generating  hydrogen.  Fill  the  generator  to  the 
depth  of  about  1  cm.  with  copper  clippings  or  cut  pieces 
of  sheet  copper.  In  another  vessel,  dilute  some  nitric  acid 
with  twice  its  volume  of  water.  Cool  this  liquid  and  pour 
it  upon  the  copper  until  the  generator  is  one-third  filled. 
Collect  the  gas  over  water.  It  is  well  to  stand  the  gener- 
ator in  cold  water.  If  the  evolution  of  gas  becomes  too 
energetic,  pour  a  little  water  down  the  funnel-tube.  Col- 
lect four  bottles  and  one  test-tube  of  the  gas.  Then  stop, 
as  even  the  first  gas  which  comes  off  is  not  quite  pure  and 
the  later  portions  contain  large  quantities  of  nitrous  oxide 
and  other  impurities.  Leave  the  bottles  standing  in  the 
water. 

Take  the  apparatus  to  the  hood,  disconnect  it  (do  not 
inhale  the  gas)  and  pour  a  little  of  the  blue  liquid  into  a 
clean  dish.  Put  a  small  flame  under  it  and  let  it  evaporate 
almost  to  dryness  under  the  hood.  What  is  the  product? 
Wash  off  the  copper  in  the  generator  and  return  what  is 
left  of  it  to  the  stock-bottle. 

Expose  a  bottle  of  nitric  oxide  to  the  air.  (?)  Into 
another  bottle  slowly  admit  oxygen  from  a  cylinder  or  gas 
holder.  Stop  after  adding  a  few  bubbles  of  oxygen,  and 
shake  the  bottle,  keeping  its  mouth  under  water ;  then  con- 
tinue adding  oxygen.  Explain  these  results. 

Slip  aside  the  cover  of  a  third  bottle  and  plunge  to  the 
bottom  of  the  bottle  some  sulphur,  burning  vigorously  in 
a  spoon.  Eemove  the  sulphur  at  once,  keeping  the  bottle 
covered,  and  do  the  same  thing  with  a  lighted  candle. 
Why  should  the  candle  be  placed  in  the  bottom  of  the  bot- 
tle ?  Interpret  the  results. 


78  ELEMENTARY  CHEMISTRY 

Into  the  fourth  bottle  introduce  a  piece  of  burning 
magnesium  held  in  forceps.  (?)  Which  flame  has  the 
higher  temperature,  that  of  a  candle  or  magnesium? 
Which  gas  requires  the  higher  temperature  to  decompose 
it,  nitrous  or  nitric  oxide?  Construct  a  tabular  view  of 
the  properties  of  the  two  gases. 

EXPERIMENT  59. — Nitrogen  peroxide. — Place  in  a  dry 
sealed  tube  of  hard  glass  lead  nitrate  to  the  depth  of  about 
1  cm.  and  heat  gently.  The  crackling  which  often,  occurs  is 
due  to  the  fact  that  the  crystals  inclose  small  drops  of 
water  and  fly  to  pieces  when  this  water  is  vaporized  by 
heat.  What  two  gases  are  given  off  ?  What  is  left  in  the 
tube?  Bring  into  the  mouth  of  the  tube  a  piece  of  moist 
blue  litmus  paper.  (?)  The  result  is  due  to  the  fact  that 
nitrogen  peroxide  reacts  with  water,  producing  nitric  acid. 

EXPERIMENT  60. — Formation  of  nitric  acid  from  water 
and  air  under  the  influence  of  electric  discharges. — Select 
a  small  glass  flask  or  bottle  and  a  2-hole  rubber  stopper 
which  fits  it.  The  holes  in  the  stopper  should  be  small. 
Pass  through  each  hole  a  platinum  wire  and  jam  a  piece 
of  glass  rod  into  each  hole  to  fix  the  position  of  the  wires. 
Both  wires  should  reach  nearly  to  the  bottom  of  the  bottle, 
and  should  be  bent  toward  each  other  so  that  the  ends  are 
less  than  1  cm.  apart.  They  must  not  touch.  The  bottom 
of  the  bottle  must  be  wet.  Throw  a  piece  of  moist  blue 
litmus  paper  into  the  bottle,  insert  tightly  the  cork  with 
the  wires,  and  pass  a  stream  of  sparks  from  a  coil  for 
half  an  hour,  or  until  the  litmus  paper  changes  color. 
The  result  is  due  to  the  production  of  nitric  acid.  What 
light  does  this  experiment  throw  upon  the  presence  of 
nitric  acid  in  rain  (paragraph  132,  Part  I)  ?  Upon  the 
supply  of  nitrogen  to  plants? 

EXPERIMENT  f>l. — Preparation  of  nitric  acid. — The 
gases  given  off  when  nitric  acid  acts  upon  metals  and 
the  vapor  of  nitric  acid  itself  are  poisonous  and  must 


COMPOUNDS  OF  NITROGEN  AND  OXYGEN 


79 


FIG.  22. 


not  be  inhaled.     Do  not  (jet  nitric  acid  upon  tlie  skin  or 
do  thin  y. 

Fit  up  the  apparatus  shown  in  Fig.  22.  Roughly 
weigh  on  the  platform  scales  enough  potassium  nitrate  to 
fill  the  retort  about  one-third.  Weigh  off  in  a  beaker  an 
equal  quantity  of  strong 
sulphuric  acid.  Intro- 
duce the  potassium  ni- 
trate into  the  retort  by 
means  of  a  piece  of 
paper.  Add  the  sul- 
phuric acid,  insert  the 
glass  stopper  (not  a  rub- 
ber or  wooden  cork, 
which  would  be  rapidly 
destroyed  by  the  nitric 
acid)  and  heat  gently. 
Only  vapor,  not  solid  or  liquid,  must  pass  over.  Why? 

Examine  the  nitric  acid  which  collects  and  record  its 
properties.  Use  it  in  the  following  experiments.  Cover 
a  piece  of  zinc  with  water  in  a  test-tube  and  slowly  add 
nitric  acid.  Do  the  same  thing  with  a  little  iron  filings 
and  a  piece  of  magnesium.  In  these  reactions  water  is 
produced,  together  with  the  nitrate  of  the  metal  used. 
The  gas  produced  may  be  nitrous  oxide,  nitric  oxide,  or 
nitrogen  peroxide  according  to  the  temperature  and 
strength  of  the  acid.  Do  not  try  to  write  the  equations, 
which  are  quite  difficult.  Simply  remember  the  general 
character  of  the  action  and  the  way  in  which  it  differs 
from  the  action  of  sulphuric  or  hydrochloric  acid  on  the 
same  metals. 

Drop  a  fragment  of  tin  into  strong  nitric  acid  in  a 
test-tube.  What  are  the  products  (paragraph  154, 
Part  I)  ? 

Drop  a  piece  of  cork  into  nitric  acid  and  let  it  remain 


80  ELEMENTARY  CHEMISTRY 

there  for  a  time.  Let  a  drop  of  nitric  acid  fall  upon  a 
colored  fabric  of  any  kind.  What  is  the  action  of  nitric 
acid  upon  organic  matter  ?  Would  the  formula  of  the  acid 
lead  you  to  expect  such  action  ? 

EXPERIMENT  62. — Aqua  regia. — Pick  up  a  small  piece 
of  gold  leaf  with  the  end  of  a  wet  glass  rod,  and  rinse  it 
into  a  test-tube  with  water.  In  the  same  way  place  another 
piece  of  gold-leaf  in  another  test-tube.  Add  to  one  tube 
nitric  acid  and  to  the  other  hydrochloric  acid.  Heat  both 
tubes.  Does  the  gold  dissolve.  Pour  the  contents  of  either 
tube  into  the  other,  and  go  on  heating.  (?)  The  product 
is  gold  chloride,  and  the  gold  dissolves  because  the  oxygen 
of  the  nitric  acid  liberates  chlorine  from  the  hydrochloric 
acid  and  the  chlorine  attacks  the  gold.  Suggest  another 
substance  which  could  be  used  with  the  hydrochloric  acid 
in  place  of  the  nitric  acid,  though  not  so  conveniently. 
The  mixture  of  nitric  and  hydrochloric  acids  will  also  dis- 
solve platinum. 

EXPERIMENT  63. — The  nitrates. — Examine  potassium 
nitrate  and  sodium  nitrate.  Try  the  solubility  of  each  in 
water.  Place  a  drop  of  each  solution  on  a  clean  glass 
plate,  let  it  evaporate  to  dryness,  and  examine  the  crystals 
left  with  a  good  lens.  Record  your  observations. 

In  a  small  iron  or  nickel  dish  melt  enough  potassium 
nitrate  to  half  fill  it  and  heat  the  melted  salt  with  the 
full  power  of  the  burner  flame.  Be  careful  not  to  upset 
the  dish,  as  the  melted  substance  would  burn  deeply  into 
the  table.  Throw  into  it  a  piece  of  charcoal  which  has 
been  heated  at  one  corner  to  redness.  Stand  aside  as  the 
melted  substance  may  sputter.  Throw  a  fragment  of  sul- 
phur into  the  dish.  Is  the  oxygen  of  potassium  nitrate 
firmly  or  loosely  held?  Does  the  result  throw  any  light 
upon  the  behavior  of  the  potassium  nitrate  in  gunpowder  ? 
Weigh  out  6  grams  of  powdered  potassium  nitrate,  1  gram 
of  powdered  charcoal,  and  1  gram  of  flowers  of  sulphur. 


COMPOUNDS  OP  NITROGEN  AND  OXYGEN          81 

Mix  thoroughly  by  pouring  from  one  paper  to  another — 
NOT  IN  A  MORTAR.  Place  the  heap  upon  a  piece  of  wood, 
thrust  a  piece  of  filter-paper  into  it,  light  the  paper  and 
stand  aside.  If  the  paper  fails  to  ignite  the  mixture, 
wait  until  the  last  spark  of  the  burning  paper  is  extin- 
guished before  approaching  the  heap  to  make  a  second 
attempt.  Explain. 

Place  a  crystal  of  silver  nitrate  on  charcoal  and  let 
the  burner  flame  play  upon  it.  (?) 

EXPERIMENT  64. — The  nitrites. — Examine  sodium  ni- 
trite. Dissolve  some  of  it  in  water.  Add  to  a  portion  of 
the  solution  a  drop  or  two  of  silver  nitrate  solution.  To 
another  portion  add  carefully  a  little  sulphuric  acid.  In- 
terpret the  result  after  reading  paragraph  159,  p.  117, 
Part  I. 

QUESTIONS 

1.  What  is  the  explanation  of  the  peculiar  conduct  of  nitrous 
and  nitric  oxides  toward  combustible  substances  ? 

2.  Vitali,  an   Italian   experimenter,  plunged  cats  into  nitric 
oxide,  and  observed  that  they  died  in  convulsions.     Why  is  it 
that  this  experiment  does  not  furnish  any  information  about  the 
action  of  nitric  oxide  upon  the  animal  body  ? 

3.  Sodium  nitrate  is  much  cheaper  than  the  potassium  salt ; 
but  supposing  that  the  two  sold  at  the  same  price,  why  would  it 
be  more  profitable  to  use  the  sodium  compound  in  making  nitric 
acid? 

4.  Why  do   railroads  object  to  transporting  nitric  acid  if 
packed  in  glass  vessels  ? 

5.  Describe  in  general  terms  the  action  of  nitric  acid  upon 
metals,  and  explain  what  becomes  of  the  hydrogen  of  the  acid. 

6.  Nitrous  oxide  yields  the  spark  test  and  supports  combus- 
tion brilliantly.     How  could  it  be  distinguished  from  oxygen  ? 

7.  Nitrogen  and  chlorine  do  not  combine  directly  to  produce 
nitrogen  chloride  ;  but  supposing  they  did,  would  heat  or  cold  be 
produced  ?     Why  ? 

8.  Why  is  it  impossible  to  employ  nitrogen  chloride  practically 
as  an  explosive  ? 


82  ELEMENTARY  CHEMISTRY 

PKOBLEMS 

54.  What  weight  of  nitric  acid  containing  80  per  cent  HNO3 
is  necessary  to  dissolve  10  grams  of  cupric  oxide  ? 

CuO  +  2HN08  =  Cu(N08)a  +  H2O. 

55.  What  weight  of  pure  nitric  acid  would  contain  50  grams 
of  oxygen  ? 

56.  Assuming  that  the  density  of  pure  nitric  acid  is  \  .5,  how 
mucli  oxygen  do  3  liters  of  it  contain  ? 

57.  How  much  nitric  acid  can  be  obtained  (a)  by  heating  200 
kilos  of  sodium  nitrate  with  sulphuric  acid;  (&)  by  heating  200 
kilos  of  potassium  nitrate  with  sulphuric  acid  ? 

58.  I  require  120  grams  of  cupric  oxide.     How  much  crystal- 
lized cupric  nitrate  must  be  heated  to  redness  to  make  it  ? 

Cu(N03)a3HaO  =  CuO  +  2N02  +  O  +  3H2O. 

59.  How  much  nitric  acid  is  needed  to  convert  400  grams  of 
potassium  hydroxide  into  potassium  nitrate  ? 

KOH  -|-  HN03  =  KN03  +  H2O. 

60.  Calculate  the  percentage  composition  of  N2Os. 

61.  Calculate  the  formula  of  a  compound  having  the  following 
composition : 

9.09  per  cent  nitrogen; 
20.77       "        oxygen; 
70.13       "         silver. 
What  is  the  name  of  the  compound  and  how  could  you  make  it  ? 


CHAPTER    XVI 

ATOMIC  AND  MOLECULAR  WEIGHTS-AVOGADR&S  RULE 

No  experiments. 

QUESTIONS 

Use  the  utmost  precision  in  answering  these  questions. 
The  subject  is  very  important. 

1.  Explain  exactly  what  is  meant  by  the  term  molecular  weight. 

2.  What  is  the  exact  difference  of  meaning  between  the  follow- 


ATOMIC  AND  MOLECULAR   WEIGHTS  83 

ing  two  formulas—  a.  NO2  ;  &.  N2O4  ?  Answer  this  question  first  in 
the  language  of  grams  and  liters.  Then  answer  again  in  the  lan- 
guage of  the  atomic  theory. 

3.  Why  is  the  density  of  a  gas  or  vapor  referred  to  hydrogen 
equal  to  one  half  the  molecular  weight  ?     Answer  this  question 
first  in  the  language  of  grams  and  liters  ;  then  in  the  language  of 
the  atomic  theory. 

4.  Explain  how  to  calculate  from  the  formula  the  weight  of 
one  liter  of  any  gas  or  vapor. 

5.  Explain  how  to  calculate  from  the  formula  the  volume  of 
one  gram  of  any  gas  or  vapor.     For  what  temperature  and  pressure 
are  the  results  obtained  in  4  and  5  good  ? 

6.  Show  that  18  grams  of  water,  HaO, 

58.5  "       "  salt,  NaCl, 

63      "       "  nitric  acid,  HNOS, 
2      "       "  hydrogen,    H,, 

must  all  contain  the  same  number  of  molecules. 

7.  Explain  Avogadro's  Rule. 

8.  What  is  the  law  of  simple  volume  ratios  ?    Why  would  the 
atomic  theory  lead  us  to  expect  it  to  be  true  ? 

9.  What  basis  of  fact  is  required  in  order  to  double  a  formula  ? 

PROBLEMS 

62.  Calculate  the  molecular  weights  of  the  following  com- 

pounds : 

a.  Sugar,  CiaHaaOn  ; 

I.  Bismuth  nitrate,  Bi(NO8)»  ; 

c.  Nitroglycerin,       C3H6(NO8),; 

d.  Glucose,  CeHiaO.. 

63.  What  is  the  weight  of  28  liters  (a)  of  nitrous  oxide  ?  (6)  of 
nitric  oxide  ? 

64.  What  is  the  volume  (a)  of  11  grams  of  nitrous  oxide  ?  (&) 
of  5  grams  of  nitric  oxide  ? 

65.  What  volume  of  hydrogen  is  needed  to  form  water  with 
the  oxygen  (a)  of  22  liters  of  nitrous  oxide  ?  (5)  of  22  grams  of 
nitrous  oxide  ? 


66.  What  is  the  volume  of  13  grams  of  nitric  oxide  ? 

67.  What  volume  of  nitrous  oxide  —  measured  at  standard  con- 

28 


84  ELEMENTARY  CHEMISTRY 

ditions — can  be  made  from  80  grams  of  ammonium  nitrate  ?    Solve 
by  inspection. 

68.  How  much  ammonium  nitrate  is  needed  to  make  80  liters 
of  nitrous  oxide  ? 

69.  How  much  ammonium  nitrate  is  needed  to  make  4,000  c.c. 
of  nitrous  oxide  ? 

70.  How  much  copper  is  needed  to  produce  30  liters  of  nitric 
oxide  ? 

71.  How  much  ammonium  nitrate  is  needed  to  produce  10 
liters  of  nitrous  oxide  ? 

72.  What  volume  of  nitrous  oxide  at  15°  C.  and  700  mm.  can 
be  obtained  by  heating  200  grams  of  ammonium  nitrate  ? 

73.  What  volume  of  nitric  oxide  at  13°  and  740  mm.  is  ob- 
tained by  dissolving  80  grams  of  copper  in  nitric  acid  ? 


CHAPTER   XVII 

ACIDS,  BASES,  AND  SALTS-METALS  AND  NON-METALS 

EXPERIMENT  65. — The  reaction  of  sodium  hydroxide 
with  hydrochloric  acid. — Dissolve  a  piece  of  sodium  hy- 
droxide about  5  cm.  (2  inches)  long  in  10  c.c.  of  water  in 
a  beaker.  Carefully  and  slowly  add  strong  hydrochloric 
acid,  stirring  constantly.  Is  there  evidence  of  energetic 
action  ?  A  white  solid  separates.  What  is  it  ?  What  else 
was  formed  at  the  same  time? 

Let  the  solid  settle  and  pour  off  the  liquid.  Dissolve 
the  solid  in  the  smallest  possible  quantity  of  water,  trans- 
fer it  to  a  dish,  and  evaporate  it  slowly  to  dryness.  When 
the  residue  is  perfectly  dry,  let  it  cool  and  examine  it. 
Taste  it.  (?)  Write  the  equation. 

EXPERIMENT  66. — The  reaction  of  sodium  hydroxide 
with  sulphuric  acid. — Make  a  little  very  strong  solution 
of  sodium  hydroxide  in  a  test-tube.  Place  about  2  c.c.  of 
the  liquid  in  another  test-tube  and  clamp  it  vertically. 
Drop  strong  sulphuric  acid  into  it — one  drop  at  a  time — • 


ACIDS,   BASES,  AND  SALTS  85 

from  a  tube  drawn  out  to  a  jet.  Be  careful.  The  reaction 
is  violent.  What  are  the  two  products?  Is  heat  evolved 
or  absorbed  ?  Write  the  equation. 

An  acid  and  a  base  always  react,  forming  a  salt  and 
water  when  they  are  brought  together,  but  the  reaction 
is  not  always  as  energetic  as  in  this  instance. 

EXPERIMENT  67. — Properties  of  acids  and  bases. — Pre- 
pare very  dilute  hydrochloric,  nitric,  sulphuric,  and  acetic 
acids  by  diluting  the  acids  with  about  100  times  their 
volume  of  water.  Make  at  least  300  c.c.  of  the  hydro- 
chloric and  nitric:  a  small  quantity  only  of  the  others. 
Taste  each  of  these  liquids.  Put  a  drop  on  red  litmus 
paper.  On  blue  litmus  paper.  On  turmeric  paper. 

Dilute  some  ammonia  with  about  100  volumes 
of  water.  Dissolve  about  3  grams  of  sodium  hy- 
droxide in  300  c.c.  of  water.  Dissolve  3  grams  of 
potassium  hydroxide  in  300  c.c.  of  water.  Obtain 
some  clear  lime-water — this  is  already  sufficiently 
dilute.  Taste  these  liquids,  rub  them  between  the 
fingers,  and  test  them  with  the  same  papers  which 
were  employed  for  the  acids. 

EXPERIMENT   68.  —  Neutralization. —  (1) 
Clamp  two  burettes  (Fig.  23)  vertically  side  by 
side.      Fill    one    with    dilute    hydrochloric    acid 
(from  Experiment  67),  and  let  the  liquid  slowly 
drop  out  until  the  bottom  of  the  meniscus  (Fig. 
7,  Part  II)   is  at  0.     Fill  the  other  to  0  with 
sodium    hydroxide    solution    (from    Experiment 
67).     Measure  out  10  c.c.  of  hydrochloric  acid 
into  a  clean  beaker,  and  add  a  few  drops  of  litmus    Flo~  553 
solution — enough    to    color    the    liquid    faintly. 
Now  allow  sodium  hydroxide  solution  to   run  into  the 
liquid,  stirring  with  a  glass  rod  until  a  change  in  color 
occurs.     This  liquid  must  fall  one  drop  at  a  time,  and 
the  change  in  color  should  be  produced  by  a  single  drop. 


86  ELEMENTARY  CHEMISTRY 

Before  this  last  drop  was  added  the  liquid  was  acid — that 

i 

is,  it  contained  H  ions;  afterward  it  was  alkaline,  it  con- 
tained OH  ions.  Read  off  the  level  of  the  liquid  in  the 
burette,  and  calculate  how  much  of  your  sodium  hydroxide 
would  be  required  for  1  c.c.  of  your  acid. 

Now  take  20  c.c.  of  the  hydrochloric  acid  and  add, 
instead  of  litmus,  2  or  3  drops  of  a  solution  of  phenol 
phthalein.  The  liquid  will  remain  colorless.  Stir  the 
liquid  constantly  and  find  out  just  how  much  of  the  so- 
dium hydroxide  solution  is  required  to  make  it  alkaline. 

In  this  case  the  OH  ions,  when  there  are  any  present,  will 
produce  a  red  color. 

Take  the  reading  and  calculate  again  how  much  of  the 
sodium  hydroxide  is  required  for  1  c.c.  of  the  acid.  Do 
the  two  results  agree?  Repeat  until  you  are  sure  that 
there  is  a  fixed  relation  between  the  quantities  of  the  two 
liquids  which  react. 

What  becomes  of  the  OH  ions  which  are  added  before 
the  color  change  occurs? 

(2)  Repeat,  using  potassium  hydroxide  and  nitric  acid 
— the  dilute  liquids  from  Experiment  67.  This  time  meas- 
ure off  10  c.c.  of  the  base  and  add  the  acid  drop  by  drop 
from  its  burette.  Use  litmus  as  the  indicator.  Take  20 
c.c.  of  the  base  the  second  time  and  use  a  few  drops  of  a 
solution  of  cochineal,  easily  made  by  digesting  the  crushed 
insects  in  alcohol  for  some  hours.  It  is  orange  when  acid 
and  violet  when  alkaline. 

If  time  permits,  repeat  ( 1 ) ,  using  potassium  hydroxide 
instead  of  sodium  hydroxide  with  the  hydrochloric  acid. 
Calculate  the  volume  of  your  potassium  hydroxide  which 
would  be  equivalent  in  neutralizing  power  to  1  c.c.  of  your 
sodium  hydroxide.  Thus  suppose  that 


ACIDS,   BASES,  AND  SALTS  87 

8  c.c.  NaOH  are  equivalent  to  10  c.c.  HC1, 
and     12  "    KOH      "          "  "  10  "       " 

clearly  8  "    NaOH    "          "  "  12  "    KOH, 

and       1  "    ISTaOH  is  "  "  J^,  or  1.5  c.c.  KOH. 

Repeat  (2),  using  sodium  hydroxide  instead  of  potas- 
sium hydroxide  with  the  nitric  acid.  Calculate  again  the 
volume  of  your  KOH  solution,  which  is  equivalent  to  1 
c.c.  of  your  NaOH.  Does  the  result  agree  with  that  ob- 
tained above  ?  If  not,  repeat.  Since  the  burette  only  reads 
to  0.1  c.c.,  do  not  look  for  absolute  agreement. 

EXPERIMENT  69.  —  Electrolysis.  —  Place  some  dilute 
solution  of  copper  sulphate,  CuS04,  in  a  beaker.  In  the 
liquid  place  two  pieces  of  platinum  foil  each  about  2  cm. 
square.  Each  is  attached  to  an  insulated  copper  wire 
which  leads  to  the  terminal  of  an  Edison-Lalande  cell,  or 
some  other  source  of  the  electric  current.  It  is  well  to 
have  a  cheap  galvanometer  in  the  circuit.  Pass  the  cur- 
rent for  fifteen  minutes  and  observe  the  result.  The  cur- 
rent evidently  passes  through  the  liquid.  What  carries 
it?  Does  the  copper  go  to  the  negative  or  positive  pole? 
Why?  Reverse  the  direction  of  the  current  for  a  time.  (?) 
The  platinum  foil  can  be  cleaned  with  a  little  nitric 
acid. 

Dissolve  some  sugar  in  about  twenty  times  its  weight 
of  distilled  water,  and  introduce  the  electrodes.  A  galva- 
nometer must  be  in  the  circuit.  Does  the  sugar  solution 

conduct?     Sugar  contains  much  hydrogen.     Its  formula 

+ 
is  C^H^On.    Is  it  an  acid?    Does  it  yield  H  ions  when 

dissolved  in  water  ?    Verify  your  conclusion  by  testing  the 
solution  with  red  and  blue  litmus  paper. 

Dilute  some  alcohol  with  about  ten  times  its  volume 
of  water  and  introduce  the  electrodes.  Does  it  conduct? 
Alcohol  is  the  hydroxide  of  a  radical.  Its  formula  is 

C2H6OH.    Is  it  a  base?    Does  it  dissociate,  yielding  OH 


88  ELEMENTARY   CHEMISTRY 

ions  when  mixed  with  water?  Verify  your  conclusion  by 
testing  the  liquid  with  both  kinds  of  litmus  paper. 

Add  very  dilute  sulphuric  acid  (about  1  to  100  parts 
of  water)  to  some  blue  litmus  solution,  one  drop  at  a 
time,  stirring  constantly  until  the  color  is  purple,  half- 
way between  blue  and  red.  If  you  put  in  too  much  acid, 
making  the  color  red,  add  very  dilute  ammonia.  Dissolve 
some  sodium  sulphate  in  water,  color  the  liquid  with  this 
neutral  litmus,  place  the  liquid  in  a  U-tube,  and  pass  the 
current  for  a  time.  Is  any  metal  deposited  at  the  negative 
pole?  What  is  formed  there?  Salts  of  the  metals  of  the 
sodium  and  calcium  groups  always  behave  in  this  way 
when  the  current  is  passed  through  water  solutions.  What 
happens  at  the  positive  pole? 

EXPERIMENT  70. — Metals  and  non-metals, — As  a  typ- 
ical solid  non-metal,  examine  sulphur.  As  a  typical  metal, 
examine  lead.  Notice  the  difference  in  luster.  If  the  lead 
is  tarnished,  scrape  it  to  expose  a  fresh  surface.  Rub  the 
sulphur  vigorously  on  the  coat-sleeve,  and  bring  it  near 
to  some  fragments  of  paper.  Notice  that  it  receives  a 
charge  and  retains  it.  This  shows  that  the  sulphur  is  a 
very  bad  conductor  of  electricity.  Treat  the  lead  in  a  sim- 
ilar way,  and  notice  that  the  charge  is  immediately  con- 
ducted away  through  the  lead  and  the  body  to  the  earth. 
The  lead  conducts  the  current.  The  same  difference  exists 
between  the  capacity  of  the  two  substances  for  conducting 
heat. 

Investigate  the  tenacity  of  the  two  by  endeavoring  to 
pull  apart  a  fragment  with  the  hands.  (?)  Investigate 
their  behavior  under  the  hammer.  What  are  the  chief 
physical  differences  between  metals  and  non-metals? 
What  are  the  chemical  differences? 

As  an  example  of  an  element  on  the  border-line  be- 
tween the  two  classes  examine  antimony.  Has  it  a  metallic 
luster?  Is  it  malleable? 


ACIDS.  BASES,  AND  SALTS  89 

QUESTIONS 

1.  What  is  an  acid  ?    What  is  a  base  ?    What  happens  when 
the  two  are  brought  together  ?     What  is  a  salt  ? 

2.  How  do  acids  act  upon  oxides  ?     Upon  carbonates  ?     Why 
is  it  that  no  gas  escapes  when  an  acid  acts  upon  an  oxide  or  a 
hydroxide  ? 

3.  What  is  an  ion  ?    Give  two  examples  of  ions  which  are 
single  atoms.     Give  two  examples  of  ions  which  are  groups  of 
atoms. 

4.  What  is  the  exact  difference  between  chlorine  ions  and 
chlorine  gas  ? 

5.  Sodium,  like  other  metals,  contains  but  a  single  atom  in  its 
molecule.     What,  then,  is  the  difference  between  metallic  sodium 
and  sodium  ions  ? 

6.  According  to  the  idea  of  electrolytic  dissociation,  when  the 
substances  are  all  dissolved  in  water, 

+  — 

Sodium  hydroxide  consists  of  Na  and  OH  ; 

Hydrochloric  acid       "        "  H  and  Cl ; 
Sodium  chloride          "        "  Na  and  Cl. 

Between  what  two  ions  does  the  actual  reaction  occur  when 
NaOH  and  HC1  are  brought  together  ?  Discuss  exactly  what  hap- 
pens to  all  four  ions. 

7.  Using  the  dissociation-idea,  explain  the  fact  that  all  salts 
of  the  same  metal  when  dissolved  in  abundant  water  produce  the 
same  color. 

8.  Is  electrolytic  dissociation  a  fact  or  a  supposition  ? 

PROBLEMS 

74.  10  grams  of  pure  sodium  hydroxide  are  dissolved  in  water. 
a.  How  much  nitric  acid  must  be  added  to  make  the  solution  neu- 
tral ?     5.  How  much  sodium  nitrate  would  be  obtained  if  this  was 
done  ? 

75.  I  have  a  solution  which  contains  just  40  grams  of  pure 
sodium  hydroxide  in  1  liter.     Calculate  the  quantities  by  weight 
of  (a)  HC1,  (K]  HNOs,  and  (c)  H3SO4,  which  will  be  required  to 
neutralize  1  c.c.  of  it. 

76.  15.75  grams  of  nitric  acid  are  mixed  with  23.25  grams  of 


90  ELEMENTARY  CHEMISTRY 

sodium  hydroxide,  both  dissolved  in  water.    What  two  compounds 
does  the  solution  contain  and  how  much  ? 

77.  In  ascertaining  the  strength  of  a  dilute  solution  of  HC1,  50 
c.c.  of  it  were  measured  out  and  neutralized  with  a  solution  of 
sodium  hydroxide  containing  .003  gram  of  NaOH  in  1  c.c.     40 
c.c.  of  the  sodium  hydroxide  solution  was  required.    What  weight 
of  HC1  was  contained  in  1  c.c.  of  the  hydrochloric  acid  ? 

78.  30  c.c.  of  a  solution  of  potassium  hydroxide  containing  .01 
gram  of  KOH  in  1  c.c.  was  needed  to  neutralize  40  c.c.  of  a  solu- 
tion of  HC1.     How  much  HC1  did  15  c.c.  of  the  hydrochloric-acid 
solution  contain  ? 

79.  20  c.c.  of  a  solution  containing  .005  gram  of  KOH  in  1  c.c. 
just  neutralized  20  c.c.  of  a  solution  of  hydrochloric  acid.    How 
much  HC1  did  15  c.c.  of  the  latter  contain  ? 


CHAPTER   XVIII 

THE  SODIUM   GROUP 

EXPERIMENT  71. — Potassium. — Potassium  is  preserved 
under  naphtha,  and  the  bottle  containing  it  must  be  kept 
stoppered  and  not  opened  in  the  vicinity  of  a  flame.  It 
catches  fire  on  contact  with  water,  and  all  apparatus  used 
in  handling  it,  as  well  as  the  desk,  must  be  dry.  It  must 
not  be  touched  with  the  fingers.  When  exposed  to  the  air 
it  takes  fire  spontaneously  after  a  time;  therefore  it  must 
not  be  put  away  under  the  desk  or  left  lying  around.  When 
thrown  into  water  it  reacts  explosively,  and  a  glass  plate 
should  be  used  to  protect  the  eyes  when  the  experiment  is 
performed. 

Examine  a  piece  of  potassium.  Cut  it  and  notice  tho 
luster.  Is  it  permanent?  Why?  Add  2  or  3  drops  of 
a  solution  of  phenol  phthalein  to  a  bottle  half  full  of 
water,  and  throw  in  a  piece  of  potassium  half  the  size  of 
a  small  pea  (no  larger).  What  happens?  What  sub- 


impo 
/x-  "E 


THE  SODIUM  GROUP  91 

stances  are  formed?  Does  the  liquid  contain  hydroxyl 
ions? 

Scrape  a  hollow  in  a  fragment  of  charcoal,  place  in  it 
a  fragment  of  potassium  half  the  size  of  a  pea,  and  let 
the  burner  flame  play  upon  it.  Notice  the  flame  color. 

Make  in  the  form  of  a  small  table  a  comparison  be- 
tween potassium  and  sodium.  Include  in  this  comparison 
luster,  permanence  of  luster,  hardness,  flame  color,  beha- 
vior with  water,  and  any  other  properties  that  may  appear 
important. 

EXPERIMENT  72. — Potassium  compounds. — Examine 
potassium  hydroxide,  chloride,  chlorate,  bromide,  and  car- 
bonate. Test  each  for  solubility  in  water.  Make  the 
flame  test  with  each,  using  a  clean  iron  wire,  as  with  so- 
dium. Clean  the  wire  after  each  test  by  dipping  it  in  a 
little  hydrochloric  acid  in  a  beaker  and  holding  it  in  the 
flame. 

Make  a  small  quantity  of  a  mixture  of  potassium  and 
sodium  chlorides  and  apply  the  flame  test  to  the  mixture. 
Look  at  the  flame  of  the  mixture  through  blue  glass. 

Make  the  flame  test  with  lithium  chloride,  LiCl.  If  a 
spectroscope  is  available,  use  it  to  observe  the  spectra  of 
potassium,  sodium,  and  lithium  chlorides  separately. 

QUESTIONS 

1.  What  is  the  connection  of  potassium  with  plant  life  ? 

2.  From  the  standpoint  of  electrolytic  dissociation,  what  really 
happens  when  potassium  acts  upon  water  is  that  the  potassium 
passes  from  one  condition  to  another.     Explain  this  statement. 

3.  From  the  same  standpoint,  explain  why,  if  a  metal  acts  vio- 
lently upon  water,  we  should  expect  its  hydroxide  to  be  a  strong 
base  and  its  salts  to  be  good  conductors  of  the  electric  current. 

PROBLEMS 

80.  How  much  potassium  is  required  to  liberate  from  water 
enough  hydrogen  to  combine  with  3  grams  of  oxygen  ? 


92  ELEMENTARY   CHEMISTRY 

81.  What  volume  of  oxygen  is  needed  to  combine  with  the 
hydrogen  given  off  by  the  action  of  9.75  grams  of  potassium  on 
water  ? 

82.  How  much  (a)  potassium  carbonate  must  be  heated  with 
how  much  (b)  pure  charcoal  to  produce  9.75  grams  of  potassium, 
and  (c)  what  volume  of  carbon  monoxide  would  be  liberated  ? 

KaCO,  +  20  =  3CO  +  2K. 


CHAPTER    XIX 

THE   COPPER    GROUP 

EXPERIMENT  73. — Copper. — Examine  copper  filings 
and  sheet  copper.  Record  its  properties.  Heat  a  piece 
of  sheet  copper  in  the  Bunsen  flame.  Place  a  little  copper 
filings  in  a  test-tube,  cover  with  water,  and  add  hydro- 
chloric acid.  Is  there  any  result?  Repeat,  using  water 
and  sulphuric  acid.  Repeat,  using  water  and  nitric  acid. 
What  is  the  gas  given  off  when  nitric  acid  dissolves  copper, 
and  why  does  it  turn  red  on  entering  the  air? 

Make  a  spiral,  about  15  cm.  long,  of  copper  wire  by 
winding  it  around  a  glass  tube.  Have  ready  a  wide  test- 
tube  and  a  cork  to  fit  it.  The  test-tube  should  contain 
about  1  c.c.  of  alcohol.  Hold  the  spiral  in  forceps  and 
pass  it  through  the  burner  flame  until  it  turns  black.  ( ?) 
Remove  it  from  the  flame,  place  in  the  test-tube,  and 
loosely  cork  the  latter.  When  cool,  remove  it  from  the 
tube  and  examine  it.  It  now  has  the  true  cok*r  of  copper. 
Explain  what  has  happened. 

EXPERIMENT  74. — Cupric  oxide. — Examine  cupric  ox- 
ide. .Mix  3  grams  of  it  with  0.5  gram  of  powdered  char- 
coal and  heat  the  mixture  in  a  hard  glass  test-tube  with 
a  delivery  tube  so  arranged  that  any  gas  given  off  must 
pass  through  lime-water  in  a  wide  test-tube.  Heat  gently 
at  first  and  then  more  intensely.  When  the  gas  has  been 


THE  COPPER  GROUP  93 

bubbling  through  the  lime-water  for  a  time,  introduce  a 
lighted  match  into  the  upper  part  of  the  wide  test-tube. 
What  gas  is  given  off?  What  must  have  happened  to  the 
copper  oxide? 

Cover  the  hard  glass  tube  with  a  layer  of  soot  from  the 
luminous  flame  and  let  it  cool  completely.  Examine  the 
contents  with  a  lens. 

How  does  carbon  affect  the  oxides  of  most  metals? 
What  is  the  importance  of  the  reaction? 

EXPERIMENT  75. — The  effect  of  cupric  oxide  upon  or- 
ganic compounds. — Heat  to  redness  in  a  hard  glass  test- 
tube,  5  grams  of  a  mixture  of  cupric  oxide  with  -fa  of  its 
weight  of  sugar,  both  finely  powdered  and  carefully  mixed. 
Conduct  the  gases  given  off  into  the  bottom  of  an  empty 
test-tube  which  stands  in  a  small  beaker  containing  a 
freezing  mixture.  Use  a  mixture  of  crushed  ice  and  salt, 
or  else  crystallized  sodium  sulphate  with  twice  its  weight 
of  hydrochloric  acid.  This  test-tube  is  closed  with  a 
doubly  perforated  cork,  and  from  it  the  gases  are  led  into 
a  second  test-tube  containing  clear  lime-water.  This  tube 
need  not  be  cooled  and  must  not  be  corked.  What  is  the 
result?  What  does  the  experiment  teach  about  the  com- 
position of  sugar?  What  becomes  of  the  copper  oxide? 
Cover  the  tube  with  soot  and  disconnect  before  cooling. 

EXPERIMENT  76. — Cupric  sulphate.  Cupric  ions. — 
Examine  cupric  sulphate.  Heat  a  little  of  it  carefully  in 
a  porcelain  dish.  (?)  Make  a  solution  of  cupric  sulphate. 
What  is  the  color  of  cupric  ions?  Place  a  portion  of  the 
solution  in  aAest-tube,  add  a  piece  of  zinc  and  a  few  drops 
of  sulphuric  acid  to  hasten  the  action.  Let  stand  for  some 
time.  The  solid  product  is  copper,  finely  divided  and  lus- 
terless.  Why  does  the  liquid  lose  its  color?  Which  has 
the  strongest  tendency  to  exist  as  ions,  copper  or  zinc? 

Kepeat  the  experiment,  using  iron  (a  nail)  in  place  of 
the  zinc. 


94:  ELEMENTARY  CHEMISTRY 

Place  a  small  piece  of  clean  sheet  copper  in  a  solution 
of  silver  nitrate.  What  is  the  product?  Which  tends 
most  strongly  to  exist  as  ion,  copper  or  silver?  What 
ought,  therefore,  to  be  the  effect  of  zinc  upon  silver  nitrate 
solution  ?  Try  it  with  a  fresh  portion. 

Add  a  drop  or  two  of  your  copper  sulphate  solution 
to  a  test-tube  half  filled  with  water,  and  investigate  the 
action  of  ammonia  water  on  the  liquid.  This  is  a  delicate 
test  for  cupric  ions. 

QUESTIONS 

1.  How  would  you  convert  cupric  ions  into  metallic  copper  ? 
Metallic  copper  into  cupric  ions  ? 

2.  What  is  the  exact  difference  between  cuprous  and  cupric 
ions  ? 

3.  The  same  electric  current  is  passed  in  succession  through  a 
solution  containing  cuprous  ions  and  one  containing  cupric  ions. 
In  one  hour  .3175  gram  of  copper  deposits  in  the  cupric  solution. 
How  much  is  deposited  in  the  cuprous  solution  in  the  same  time, 
and  why  ? 

PROBLEM 

83.  2  grams  of  finely  divided  copper  were  heated  in  oxygen. 
2.5063  grams  of  cupric  oxide  were  produced.  Calculate  the 
atomic  weight  of  copper. 


CHAPTER   XX 

SILVER 

EXPERIMENT  77. — Silver. — Dissolve  about  0.5  gram 
of  silver  nitrate  in  50  c.c.  of  distilled  water  in  a  beaker. 
Try  the  action  of  a  clean  piece  of  sheet  copper  upon  a 
portion  of  the  liquid  in  a  test-tube.  Add  a  drop  of  mer- 
cury to  another  portion.  Let  both  tubes  stand  some  time. 
Thfe  visible  product  is  silver.  What  else  must  have  been 
produced  in  both  cases? 


SILVER  95 

To  the  rest  of  the  silver  nitrate  solution  add  a  solu- 
tion of  sodium  chloride.  The  precipitate  is  silver  chlo- 
ride, AgCL  Write  the  equation.  In  what  previous  ex- 
periment did  you  make  and  study  silver  chloride?  What 
liquid  dissolves  it?  What  is  the  effect  of  light  upon  it? 
Answer  these  questions  from  your  notes.  Stir  up  the 
silver  chloride  with  water,  allow  it  to  settle,  and  pour  off 
thtvater  without  losing  any  of  the  silver  chloride.  Re- 
or  four  times.  This  is  called  washing  by  de- 
Place  a  piece  of  zinc  in  contact  with  the  silver 
chloric,  and  add  a  drop  or  two  of  sulphuric  acid.  Allow 
to  stand.  The  product  is  silver.  Wash  it  by  decantation, 
examine  it,  and  record  its  properties. 

EXPERIMENT  78. — Photography. — If  a  negative  is 
available,  examine  it.  What  is  the  material  of  which  the 
image  consists?  Describe  briefly,  from  the  beginning,  the 
processes  through  which  the  negative  has  passed.  Why 
are  the  lights  and  shadows  reversed  in  it?  If  practicable, 
expose  a  piece  of  sensitive  paper  back  of  the  negative  in  a 
printing-frame  to  sunlight.  The  paper  must  be  inserted 
by  dim  light,  and  the  sensitive  side  of  the  paper  must  be 
in  contact  with  the  side  of  the  negative  bearing  the  film. 
Once  a  minute  open  half  the  printing-frame  in  a  dim  light 
and  examine  the  progress  of  the  printing.  When  the  print 
is  finished,  remove  and  examine  it.  How  can  the  unpleas- 
ant reddish  color  be  altered?  What  is  toning,  and  how 
does  it  affect  the  chemical  composition  of  the  image?  Is 
your  picture  permanent?  Expose  it  to  strong  light  and 
find  out.  If  not,  how  could  it  be  made  permanent  ? 

EXPERIMENT  79. — Silver  oxide. — Add  some  sodium 
hydroxide  solution  to  a  dilute  solution  of  silver  nitrate. 
The  precipitate  is  silver  oxide,  Ag20.  Complete  the 
equation — 

2AgN03  +  OTaOH  =  Ag20  +...  +  ... 
What  experiment  have  you  made  upon  the  solubility  of 


96  ELEMENTARY  CHEMISTRY 

gold  ?  Is  it  soluble  in  nitric  acid  ?  In  hydrochloric  acid  ? 
In  a  mixture  of  both?  How  could  you  separate  gold 
from  silver  in  an  alloy  of  both  metals? 

QUESTIONS 

1.  How  could  you  convert  silver  nitrate  into  silver  chloride  ? 
Silver  chloride  into  silver  nitrate  ? 

2.  How  could  you  prepare  pure  silver  from  a  silver  coin  con- 
taining 90  per  cent  silver  and  10  per  cent  copper  ? 

PROBLEMS 

84.  Calculate  the  percentage  composition  (a)  of  silver  chloride, 
AgCl ;  (5)  of  silver  sulphide,  AgaS. 

85.  Calculate  the  formula  of  a  compound  of  the  following 
composition : 

Silver 65.45 

Sulphur 19. 39 

Arsenic 15. 16 

86.  When  hydrogen  is  heated  with  silver  chloride  silver  is 
produced : 

AgCl  +  H  =  HC1  +  Ag. 

If  52.65  c.c.  of  hydrogen  produce  0.505  gram  of  silver,  what  is 
the  atomic  weight  of  silver  ? 

87.  Calculate  the  percentage  composition  of  silver  acetate, 
AgC,H,O,. 

88.  How  much  zinc  is  required  to  precipitate  5  grams  of  sil- 
ver from  solution  ? 

+  ++ 

2Ag  +  Zn  =  Zn  +  2Ag. 


CHAPTER  XXI 

No  experiments. 

PROBLEMS 

89.  How  much  auric  chloride  can  be  made  from  65.67  grams 
of  gold  ? 

90.  Calculate  the  formula  of  a  compound  containing  92.8  per 
cent  of  gold  and  7.7  per  cent  of  oxygen. 


THE  CALCIUM  GROUP  97 

91.  What  volume  of  oxygen  is  produced  when  49  grams  of 
auric  oxide  are  heated  ? 


CHAPTER   XXII 

THE  CALCIUM  GROUP 

EXPERIMENT  80. — Calcium  compounds. — Support  a 
piece  of  marble  on  a  pipe-stem  triangle  and  heat  it  with 
the  flame  of  a  blast-lamp  for  fifteen  minutes.  Examine 
the  product.  Is  the  change  complete  or  partial?  What 
gas  has  escaped?  In  what  respects  would  the  result  have 
been  different  if  you  had  heated  the  marble  in  a  sealed 
vessel?  Write  the  equation. 

Place  the  mass  in  a  beaker  containing  distilled  water, 
allow  it  to  stand  for  a  time,  and  remove  the  unchanged 
marble.  Decant  the  liquid  carefully  upon  a  filter.  What 
is  the  name  and  formula  of  the  white  solid  which  is  left? 
The  liquid  which  runs  through  the  filter  is  lime-water. 
Test  its  action  upon  red  and  blue  litmus  paper.  Dissolve 
a  little  ammonium  oxalate  in  distilled  water  and  add  it 
to  some  of  the  lime-water.  The  precipitate  is  calcium 
oxalate.  This  is  a  delicate  test  for  calcium.  Test  faucet 
water  carefully  for  calcium  in  the  same  way. 

EXPERIMENT  81. — Examine  calcium  chloride.  Expose 
a  fragment  to  the  air  on  a  dry  glass  plate  for  an  hour.  (?) 
What  is  it  used  for  in  the  laboratory  ?  Make  the  flame  test 
with  it.  Make  the  flame  test  with  strontium  chloride  and 
barium  chloride.  (If  the  chlorates  are  at  hand,  use  them 
instead  of  the  chlorides.  The  result  is  more  satisfactory.) 

Make  a  very  dilute  solution  of  barium  chloride,  and 
add  to  it  a  few  drops  of  sulphuric  acid.  The  precipitate 

4  + 

is  barium  sulphate,  BaS04.    This  is  a  delicate  test  for  Ba 
ions,  and  conversely,  a  solution  of  barium  chloride  is  a 


98  ELEMENTARY  CHEMISTRY 

delicate  test  for  1304  ions ;  that  is,  for  sulphuric  acid  or 
sulphates. 

Divide  the  liquid  containing  the  precipitate  into  two 
parts,  and  show  that  the  latter  is  insoluble  in  hydrochloric 
acid  and  in  ammonia  water.  This  serves  to  distinguish 
barium  sulphate  from  other  precipitates  of  similar  ap- 
pearance. 

PROBLEMS 

92.  A  piece  of  pure  marble  weighing  10  grams  is  heated  to 
complete  decomposition,     (a)  What  is  the  formula  and  weight  of 
the  substance  which  remains  ?     (b)  What  gas  escapes,  and  what 
volume  measured  at  20°  and  740  mm.  ? 

93.  How  many  tons  of  limestone  must  be  heated  to  make  200 
tons  of  lime  ? 

94.  1.363  kilos  of  marble  are  heated  until  entirely  decom- 
posed, and  water  is  thrown  on  the  residue. 

95.  Calculate  the  percentage  composition  («)  of  barium  sul- 
phate, BaSO4;  (?>)  of  barium  carbonate,  BaCO3. 

96.  10  grams  of  barium  carbonate  are  dissolved  in  hydro- 
chloric acid,     (a)  What  volume  of  carbon  dioxide  is  produced, 
and  (Z>)  how  much  crystallized  barium  chloride  (BaCl32HaO)  can 
be  obtained  from  the  solution  ? 

97.  How  much  barium  sulphate  can  be  made  from  2  grams  of 
calcium  sulphate  (CaSO4)  ? 

98.  1.182  grams  of  barium  carbonate  were  dissolved  in  hydro- 
chloric acid,  and  the  solution  precipitated  with  sulphuric  acid. 
The  barium  sulphate  obtained  weighed  1.398  grams.     Calculate 
the  percentage  of  barium  in  the  barium  carbonate. 


CHAPTEE   XXIII 

MAGNESIUM 

EXPERIMENT  82. — Examine  magnesium.  Is  it  light 
or  heavy  ?  Is  its  luster  affected  by  the  air  ?  Hold  a  piece 
of  magnesium  ribbon  20  cm.  (8  in.)  long  in  forceps, 


MAGNESIUM  99 

and  burn  it.  Eeceive  the  product  in  a  dish  and  reserve 
it.  Burn  a  similar  piece  in  steam,  proceeding  exactly  as 
directed  on  page  173,  Part  I.  Before  introducing  the 
magnesium,  show  that  all  the  air  has  been  expelled  from 
the  beaker  by  introducing  a  burning  match  or  candle, 
which  should  be  extinguished.  Compare  the  products  of 
burning  magnesium  in  air  and  in  steam.  Are  they  iden- 
tical? Investigate  the  action  of  hydrochloric,  nitric,  and 
sulphuric  acids  separately  in  magnesium.  Cover  a  small 
piece  of  the  metal  with  water  in  a  test-tube,  and  add 
the  acid  gradually.  Record  the  results.  Remembering 
that  magnesium  is  bivalent,  write  the  equations  for  the 
action  of  hydrochloric  and  sulphuric  acids  upon  it.  The 
action  of  nitric  acid  is  more  complicated. 

EXPERIMENT  83. — Carefully  mix  on  paper  1  gram  of 
powdered  magnesium  with  1.5  grams  of  powdered  potas- 
sium chlorate.  The  materials  must  not  be  ground  to- 
gether. Place  the  mixture  on  a  brick  or  a  block  of  wood 
under  the  hood,  stick  a  piece  of  fiHer  paper  about  10  cm. 
long  in  the  heap,  light  the  end  of  the  paper  farthest  from 
the  powder,  and  step  aside.  The  action  is  explosive  and 
the  light  intense.  //  the  burning  paper  fails  to  ignite  the 
mixture,  wait  till  you  are  absolutely  sure  that  it  is  ex- 
tinguished before  approaching. 

What  are  the  products  of  the  change?  Why  is  the 
action  so  rapid?  What  application  is  made  of  mixtures 
of  magnesium  with  substances  yielding  oxygen? 

QUESTIONS 

1.  What  is  the  valence  of  magnesium  in  MgCU,  MgO,  MgSO4, 
andMg(NO,)2?     Why? 

2.  What  is  the  chemical  formula  of  calcined  magnesia?   of 
Epsom  salts  ?    Why  is  magnesium  chloride  objectionable  in  water 
for  steam  boilers  ? 

29 


100  ELEMENTARY  CHEMISTRY 

3.  Can  you  perceive  any  connection  between  the  properties  of 
magnesium  oxide  (almost  infusible  and  non-volatile)  and  the 
great  brightness  of  the  flame  of  burning  magnesium  ? 

PKOBLEMS 

99.  What  volume  of  oxygen  is  needed  to  burn  9  grams  of 
magnesium  ? 

100.  If  0.4  gram  of  magnesium  liberated  391  c.c.  of  dry  hydro- 
gen at  13°  when  treated  with  HC1,  what  is  the  atomic  weight  of 
magnesium  ? 

101.  The  electric  current  is  passed  through  fused  magnesium 
chloride  until  14  grams  of  magnesium  are  obtained.     What  volume 
of  chlorine  at  standard  conditions  is  liberated  ? 


CHAPTER    XXIV 

ZINC  AND    CADMIUM 

EXPERIMENT  84. — Examine  sheet  zinc  and  record  its 
properties.  Refer  to  your  notes  of  former  experiments 
"for  data  regarding  the  action  of  acids  upon  it.  Heat  a 
small  piece  of  zinc  upon  charcoal  with  the  blowpipe  flame. 
Practise  with  the  blowpipe  until  you  can  produce  a  con- 
tinuous flame.  The  blast  comes  from  the  cheeks,  not  from 
the  lungs,  and  the  cheeks  are  refilled  with  air  at  intervals 
without  interrupting  the  flame.  What  is  the  deposit  upon 
the  charcoal?  Repeat  Experiment  83,  using  zinc  dust 
with  the  KC103  instead  of  magnesium.  Care. 

EXPERIMENT  85.— Quantitative  experiment.  Atomic 
weight  of  zinc.— Fit  up  the  apparatus  shown  in  Fig.  24. 
T  is  a  graduated  tube  holding  100  c.c.  It  is  filled  with 
water,  inverted  in  a  vessel  of  water,  and  clamped.  The 
flask,  F9  contains  a  piece  of  pure  granulated  zinc,  which 
has  been  accurately  weighed.  It  must  not  weigh  more 
than  0.2  gram.  A  small  piece  of  platinum  wire  is  wound 


ZINC  AND   CADMIUM  101 

around   it   to   hasten   the   solution   of   the   zinc.      Before 
placing  D  under  T,  the  doubly  bored  rubber  cork  should 


be  twisted  tightly  into  the  flask,  the  end  of  D  placed  under 
water,  and  F  and  D  filled  with  water  by  pouring  water 
into  the  separating  funnel  8. 

The  stop-cock  is  then  closed  and  the  funnel  filled  with 
warm  (not  hot)  dilute  sulphuric  acid,  one  part  to  three 
of  water.  Remember  the  precautions  necessary  in  diluting 
sulphuric  acid.  The  acid  is  cautiously  allowed  to  run  into 
F,  and  more  of  it  admitted  at  intervals  until  the  zinc  is 
completely  dissolved.  The  funnel  must  never  become 
empty,  or  gas  will  escape.  Eemove  D,  cover  the  open 
end  of  T  with  the  thumb,  and  transfer  it  to  a  cylinder  of 
water  which  has  had  time  to  acquire  the  temperature  of 
the  laboratory.  Clamp  the  tube  so  that  the  level  of  water 
inside  and  out  is  the  same,  place  a  thermometer  in  the 
water,  and  let  the  apparatus  stand  15  minutes.  Read 
the  volume  of  the  gas,  the  temperature,  and  the  atmos- 
pheric pressure  on  the  barometer  in  the  laboratory.  Re- 
cord the  readings,  and  repeat  them.  Take  up  the  dif- 
ferent steps  of  the  calculation  in  the  following  order: 


102  ELEMENTARY  CHEMISTRY 

1.  Find  the  volume  your  gas  would  occupy  at  0°,  760 
mm.,  and  completely  dry.     The  method  of  making  the 
calculation  is  explained  in  the  Appendix,  and  is  the  same 
that  you  employed  in  calculating  the  weight  of  a  liter 
of  oxygen  (p.  38). 

2.  Calculate  the  weight  of  this  gas  by  multiplying  the 
corrected  volume  in  liters  by  0.0896,  the  weight  of  a  liter 
of  hydrogen  under  standard  conditions. 

3.  Calculate  by  proportion  the  weight  of  zinc  which 
would  be  required  to  set  free  2  grams  of  hydrogen.    Since 
zinc  is  bivalent,  this  will  be  the  atomic  weight  of  zinc. 

EXPERIMENT  86. — Examine  cadmium.,  and  record  its 
properties.  Has  it  the  same  luster  as  zinc?  Is  it  denser 
or  lighter?  Heat  a  chip  of  it  on  charcoal.  (?)  Dissolve 
a  little  cadmium  chloride  in  water,  and  place  a  piece  of 
zinc  in  the  liquid.  Allow  to  stand.  What  happens? 
Which  of  the  two  metals  has  the  strongest  tendency  to 
exist  as  ions? 

PKOBLEMS 

102.  If  0.5  gram  of  zinc  when  dissolved  in  hydrochloric  acid 
set  free  183.7  c.c.  of  hydrogen  measured  over  water  at  15°  and 
760  mm.,  what  is  the  atomic  weight  of  zinc  ? 

103.  If  1  gram  of  zinc  set  free  366  c.c.  of  hydrogen  measured 
over  water  at  9°  and  748  mm.,  what  is  the  atomic  weight  of  zinc  ? 


CHAPTEE   XXV 

MERCURY 

Mercury  should  not  be  brought  into  contact  with  jew- 
elry. Rings  should  be  removed  before  working  with  it. 

EXPERIMENT  87. — Examine  mercury.  Is  its  luster 
affected  by  the  air?  Notice  the  high  density  (13.6). 
Place  an  iron  nail  and  a  fragment  of  marble  upon  the 


MERC  UK  Y  103 

surface  of  mercury  in  a  beaker.     Is  there  any  other  ele- 
ment which  is  a  liquid  at  ordinary  temperatures? 

Heat  a  drop  of  it  in  a  dry,  clean  test-tube,  and  de- 
scribe the  result.  Incline  the  tube  when  introducing  the 
metal,  or  the  shock  may  break  the  tube. 

Place  a  drop  of  mercury  in  a  watch-glass  and  rub  it 
upon  a  piece  of  clean  copper.  (?)  What  is  an  amalgam? 
Heat  the  copper  carefully.  (?) 

Try  the  action  of  dilute  hydrochloric,  sulphuric,  and 
nitric  acids  upon  small  drops  of  the  metal. 

Mercury  must  not  be  thrown  into  the  sinks. 

EXPERIMENT  88.  —  Place  some  mercuric  sulphide,  HgS, 
in  a  short  tube  about  0.5  cm.  wide,  open  at  both  ends, 
and  gently  heat  the  portion  containing  the  substance,  in- 
clining the  tube  so  that  a  current  of  hot  air  shall  be 
drawn  over  the  mercuric  sulphide.  Describe  and  explain 
the  result.  Write  the  equation.  Is  this  the  usual  effect 
of  heating  with  abundant  air-supply  upon  sulphides? 
What  other  sulphides  might  be  expected  to  behave  like 
mercuric  sulphide  under  the  same  circumstances? 

EXPERIMENT  89.  —  Mercurous  compounds.  —  Add  hydro- 
chloric acid,  drop  by  drop,  to  a  solution  of  mercurous 
nitrate.  The  visible  product  is  mercurous  chloride.  Does 
it  look  like  silver  chloride?  Treat  it  with  ammonia 
water.  Does  it  behave  like  silver  chloride  with  ammonia? 

EXPERIMENT  90.  —  Mercuric  compounds.  —  Mercuric 
chloride  is  intensely  poisonous.  It  must  not  be  touched 
with  the  fingers.  Dissolve  a  little  mercuric  chloride  in 
about  50  c.c.  of  water.  To  a  portion  of  the  solution  add 
a  solution  of  sodium  hydroxide.  The  precipitate  is  mer- 
curic oxide,  HgO  —  yellow  because  it  is  finely  divided.  ^ 
Complete  the  equation— 

HgCl2  +  2NaOH  = 

To  another  portion  of  the  solution  add  a  solution  of 
potassium  iodide,  one  drop  at  a  time,  shaking  constantly. 


104  ELEMENTARY  CHEMISTRY 

The  precipitate  is  mercuric  iodide.  Notice  that  the  yellow 
modification  forms  first,  and  immediately  passes  into  the 
red  stable  modification.  Write  the  equation.  Show  that 
the  precipitate  is  soluble  in  more  potassium  iodide  solu- 
tion. 

To  the  rest  of  the  solution  add  an  equal  volume  of  a 
solution  of  stannous  chloride,  SnCl2,  and  heat  the  liquid. 
Explain. 

QUESTION 

Make  a  tabular  statement  of  the  differences  between  mer- 
curous  and  mercuric  chlorides,  using  your  experiments  and  the 
text  of  Part  I  as  sources  of  information. 

PKOBLEMS 

104.  758  grams  of  mercuric  chloride  are  dissolved  in  water. 
(a)  How  much  potassium  iodide  must  be  added  to  the  liquid,  and 
(5)  how  much  mercuric  iodide  will  be  obtained  ? 

105.  If  mercurous  chloride  contains  84.93  per  cent  mercury 
and  15.07  per  cent  chlorine,  and  if  the  formula  is  Hg3Cl3,  what  is 
the  atomic  weight  of  mercury  ? 

106.  If  mercuric  chloride  has  the  formula  HgCls  and  contains 
73.8  per  cent  of  mercury  and  26.2  per  cent  of  chlorine,  what  is  the 
atomic  weight  of  mercury  ? 

107.  88.5832  grams  of  mercuric  sulphide,  when  completely 
decomposed,  yield  76.3725  grams  of  mercury.     What  is  the  atomic 
weight  of  mercury  ? 


CHAPTEE   XXVI 

BORON  AND  ALUMINIUM 

EXPERIMENT  91. — Boric  acid  and  borax. — Dissolve 
about  10  grams  of  powdered  borax  in  50  c.c.  of  hot  water 
in  a  beaker.  Add  25-30  drops  of  strong  sulphuric  acid, 
and  allow  to  cool.  Boric  acid,  B(OH):i,  crystallizes. 

Examine  boric  acid  from  the  stock  bottle.     Hold  a 


BORON  AND  ALUMINIUM  105 

fragment  of  it  in  the  Bunsen  flame.  The  flame  color  can 
be  used  as  a  test  for  boric  acid.  Place  a  little  borax 
(0.2  gram  or  less)  in  a  dish  and  moisten  it  with  strong 
sulphuric  acid.  Stir  with  a  glass  rod,  add  5  c.c.  of  alco- 
hol, and  stir  again.  Set  fire  to  the  mixture.  Record 
the  result. 

Make  a  little  loop  in  the  end  of  a  piece  of  platinum 
wire,  heat  the  loop  red  hot,  and  dip  it  into  a  little  borax. 
Put  the  wire  with  the  borax  which  adheres  to  it  back  into 
the  flame.  What  is  the  cause  of  the  swelling  up  of  the 
borax?  What  is  the  name  and  composition  of  the  trans- 
parent bead  which  remains? 

Heat  the  bead,  and  cause  a  minute  quantity  of  man- 
ganese dioxide  to  adhere  to  it.  Melt  the  bead  again,  and 
let  it  cool.  The  color  is  characteristic  of  manganese. 
Slowly  turn  the  cap  at  the  base  of  the  burner  until  a 
small  portion  of  the  flame  is  faintly  luminous,  and  hold 
the  bead  in  this  portion  steadily.  When  cold  it  should  be 
colorless.  Heat  it  in  the  outer  portion  of  the  flame  to 
restore  the  color. 

Remove  the  bead  by  dipping  it  while  hot  into  cold 
water  and  scraping  it  off,  and  repeat  this  experiment, 
using  the  blowpipe  instead  of  the  burner  flame.  Use  a 
blowpipe-tip  on  your  burner,  closing  the  holes  at  the  base. 
The  oxidizing  flame  of  the  blowpipe  should  be  used  for 
making  the  borax  bead,  containing  the  manganese.  This 
flame  corresponds  to  the  outer  non-luminous  portion  of 
the  Bunsen-burner  flame. 

The  oxidizing  flame  is  produced  by  placing  the  end  of 
the  blowpipe  nearly  in  the  middle  of  the  flame  and  blow- 
ing steHily.  The  flame  should  be  blue,  perfectly  noiseless, 
and  parallel  to  the  slit  in  the  burner-tip. 

The  'reducing  flame  should  be  used  to  decolorize  the 
manganese  borax  bead.  To  produce  it,  place  the  tip  of 
the  blowpipe  at  the  edge  of  the  flame,  and  use  a  more 


106  ELEMENTARY  CHEMISTRY 

gentle  current  of  air  than  that  used  for  the  oxidizing 
flame.  The  flame  will  be  partly  blue  and  partly  yellow. 
It  must  be  noiseless  and  steady.  It  should  surround  the 
bead  completely,  so  as  to  prevent  the  access  of  air. 

Eemove  the  bead.  Make  a  fresh  one,  and  investigate 
the  color  produced  in  it  by  a  minute  fragment  of  copper 
sulphate  in  both  the  outer  and  inner  flames.  Record  the 
result. 

Make  another  bead,  take  up  with  it  a  scarcely  visible 
speck  of  cobalt  nitrate,  and  study  the  color  produced. 

What  use  can  be  made  of  these  phenomena  in  testing 
for  the  metals  ? 

EXPERIMENT  92. — Aluminium. — Examine  sheet  alu- 
minium. Is  it  affected  by  the  air?  How  does  its  density 
compare  with  that  of  familiar  metals,  like  iron  and  cop- 
per? How  do  small  pieces  of  it  behave  with  dilute  sul- 
phuric, nitric,  and  hydrochloric  acids?  Try  each  acid 
separately.  What  gas  is  given  off?  What  is  the  action  of 
a  solution  of  sodium  hydroxide  on  a  piece  of  aluminium? 

Examine  aluminium  powder.  This  is  the  "  bronze 
paint "  used  for  mail  boxes.  Place  a  heap  of  it  1  cm. 
wide  on  an  asbestos  plate,  and  set  fire  to  it.  What  effect 
has  aluminium  upon  metallic  oxides  at  high  temperatures, 
and  what  use  is  made  of  the  fact? 

EXPERIMENT  93.— Alum.— Dissolve  10  grams  of  alu- 
minium sulphate  in  a  little  hot  water.  In  another  beaker 
dissolve  5  grams  of  potassium  sulphate  in  a  small  quantity 
of  hot  water.  Both  liquids  must  be  clear.  If  not,  decant 
or  filter.  Mix  the  liquids  in  a  glass  dish  or  large  beaker, 
and  allow  to  cool.  The  product  is  potassium-alum.  Place 
some  of  the  crystals  on  a  glass  plate,  and  examine  them 
with  a  lens. 

QUESTIONS 

1.  What  are  some  of  the  uses  of  aluminium  ? 

2.  Why  are  there  so  many  different  alums  ?    What  chemical 


SILICON  107 

composition  must   a  substance   have  in  order  to  be   called   an 
alum  ? 

3.  Why  is  boron  considered  a  non-metal  and  aluminium  a 
metal  ? 

PKOBLEMS 

108.  Calculate  the  percentage  composition  of  borax,  Na9B4O7- 
10H3O.     Calculate  water,  not  hydrogen. 

109.  What  is  the  formula  of  a  substance  containing  31.19  per 
cent  boron  and  68.81  per  cent  oxygen  ? 

110.  If  6.75  grams  of  aluminium,  when  dissolved  in  hydro- 
chloric acid,  yield  8.4  liters  of  hydrogen,  what  is  the  atomic  weight 
of  the  metal  ? 


CHAPTEE   XXVII 

SILICON 

EXPERIMENT  94. — Hydrogen  silicide. — Double  a  piece 
of  magnesium  ribbon  6  cm.  long,  and  heat  it  in  a  sealed 
tube  of  hard  glass  to  a  bright-red  heat  for  several  min- 
utes. Allow  to  cool,  break  open  the  tube,  discard  the  mag- 
nesium if  any  remains,  and  throw  the  blackened  fragments 
of  glass  into  dilute  hydrochloric  acid,  1  part  to  2  parts 
of  water,  in  a  beaker.  The  magnesium  produces  mag- 
nesium silicide,  SiMg2,  with  the  Si02  of  the  glass.  This 
liberates  hydrogen  silicide  with  the  acid.  Write  all  the 
reactions,  and  explain  the  cause  of  the  spontaneous  igni- 
tion of  the  gas. 

EXPERIMENT  95. — Add  strong  hydrochloric  acid  to  a 
solution  of  sodium  silicate  in  a  porcelain  dish.  What  is 
the  composition  of  the  jelly  which  separates?  Evaporate 
slowly  to  dryness  under  the  hood,  allow  to  cool,  moisten 
the  residue  with  hydrochloric  acid,  half  fill  the  dish  with 
water,  heat  to  boiling,  and  filter.  The  substance  left  on 
the  filter  is  silicon  oxide,  Si02.  Wash  it  with  water  until 
a  few  drops  of  the  liquid  which  runs  through  gives  no 


108  ELEMENTARY  CHEMISTRY 

precipitate  when  collected  in  a  test-tube,  and  mixed  with 
a  drop  of  silver  nitrate  solution. 

Examine  some  crystals  of  quartz.  Make  a  drawing  of 
one  in  your  note-book.  Can  you  scratch  quartz  with  a 
knife?  With  a  file?  Is  a  quartz  crystal  hard  enough  to 
scratch  glass? 

QUESTIONS 

1.  In  what  respect  are  silicic  acid  and  nitrous  acid  similar  ? 

2.  What  is  meant  by  the  term  silicate  f 

PKOBLEMS 

111.  What  is  the  weight  of  2.8  liters  of  hydrogen  silicide 
under  standard  conditions  ? 

112.  Calculate  the  formula  of  a  silicate  which  was  found  by 
analysis  to  possess  the  following  composition  : 

Zinc 58. 6  per  cent. 

Silicon 12.7       " 

Oxygen 28.7       " 


CHAPTER   XXVIII 

TIN 

EXPERIMENT  96. — Examine  tin-foil,  bar  tin,  and  gran- 
ulated tin,  and  record  the  properties  of  the  metal.  Are 
ordinary  "tin"  vessels  made  of  solid  tin?  Investigate 
this  by  bringing  a  magnet  in  contact  with  a  tin  vessel  of 
any  kind.  Tin  is  not  attracted  by  the  magnet.  Try  this 
with  tin-foil. 

Study  the  effect  of  heat  upon  tin  by  placing  some  tin- 
foil in  a  hollow  scraped  in  a  piece  of  charcoal  and  letting 
the  blowpipe  flame  play  upon  it.  Does  it  melt?  Does  it 
absorb  oxygen?  Place  a  drop  of  mercury  in  a  watch-glass 
and  press  some  tin-foil  beneath  the  surface.  What  is  the 
product? 


TIN  109 

EXPERIMENT  97.— Heat  some  tin-foil  under  the  hood 
for  some  time  with  strong  hydrochloric  acid  in  a  wide 
test-tube.  The  result  is  a  solution  of  stannous  chloride, 
f  SnCl2.  Dilute  the  liquid,  and  place  a  piece  of  zinc  in  it. 
The  product  is  tin.  Describe  it,  and  notice  that  a  finely 
divided  metal  may  have  a  very  different  appearance  from 
the  compact  substance.  Have  you  noticed  any  other  ex- 
amples of  this  fact? 

EXPERIMENT  98. — The  atomic  weight  of  tin. — Clean 
/  carefully  a  porcelain  crucible,  dry  and  weigh  it.  Place  a 
|  little  granulated  tin  in  it  and  weigh  again  to  ascertain  the 
amount  of  tin  taken.  This  should  be  from  0.5  to  0.7 
gram.  The  quantity  must  be  accurately  known.  Cover 
the  bottom  of  the  crucible  with  water  and  slowly  add  nitric 
acid  until  the  tin  is  covered.  Support  the  crucible  on  a 
pipe-stem  triangle,  and  place  a  small  flame  5  cm.  below  the 
bottom.  This  should  be  done  under  the  hood.  When  the 
action  seems  complete,  evaporate  cautiously  to  dryness, 
avoiding  any  spattering  which  will  cause  loss  of  substance 
and  make  it  necessary  to  go  back  to  the  beginning.  When 
the  residue  is  dry,  heat  it,  gradually  at  first,  and  finally 
intensely;  let  it  cool,  and  weigh  it.  There  will  be  a  gain 
in  weight,  which  is  due  to  the  fact  that  the  tin  now  exists 
as  Sn02.  Calculate  the  atomic  weight  of  tin  thus — 

Gain  in  weight  :  02  =  weight  of  tin  :  x. 
For  example : 

Weight  of  crucible  -f  tin 16 . 642 

Weight  of  crucible,  empty 14 . 642 

Weight  of  tin 2.000 

Weight  of  crucible  -f  Sn02 17 . 180 

Weight  of  crucible  +  tin 16 . 642 

Increase  due  to  oxygen 538 

0.538  :32  =  2  :  x. 
x  =  119. 


110  ELEMENTARY  CHEMISTRY 

QUESTIONS 

1.  Discuss  the  properties  of  gray  tin  and  its  relations  to  ordi- 
nary tin. 

2.  If  20°  is  the  temperature  at  which  gray  tin  and  white  tin 
are  in  equilibrium,  what  is  the  condition  of  the  tin  of  a  roof  on  a 
winter  day  ?     How  is  it  that  tin  can  be  used  for  any  purpose 
requiring  exposure  to  cold  ? 

PROBLEM 

113.  1  gram  of  tin  was  treated  with  nitric  acid.  After  evapo- 
ration to  dry  ness  and  heating,  the  gain  in  weight  was  0.271  gram. 
Calculate  the  atomic  weight  of  tin. 


CHAPTEK   XXIX 

LEAD 

EXPERIMENT  99. — Examine  lead.  Cut  a  chip  off  a 
piece  in  order  to  observe  the  luster  of  the  untarnished 
metal.  Heat  a  little  granulated  lead  in  a  test-tube  with 
dilute  nitric  acid.  The  product  is  lead  nitrate.  Exam- 
ine litharge,  PbO,  and  heat  a  little  of  it  on  charcoal  with 
the  blowpipe  flame.  Describe  and  explain  the  result.  Ex- 
amine lead  dioxide,  Pb02.  Heat  a  little  of  it  in  a  tube 
sealed  at  one  end  (hard  glass).  Apply  the  spark  test  to 
the  gas  given  off.  What  is  the  residue?  Examine  red 
lead.  Treat  it  with  dilute  nitric  acid.  What  is  the  cause 
of  the  color-change? 

EXPERIMENT  100. — Dissolve  about  5  grams  of  lead 
nitrate  in  100  c.c.  of  water.  Use  small  portions  of  this 
liquid  to  study  the  behavior  of  lead  salts  with  the  follow- 
ing substances : 

Product. 

Sulphuric  acid,  HaSO4 Lead  sulphate,  Pb8O4 

Potassium  chromate,K2CrO4  Lead  chromate  (chrome  yellow), PbCrO4 
Hydrogen  sulphide,  HaS  . .  Lead  sulphide,  PbS 


LEAD  111 

Write  all  the  equations.  In  the  remainder  of  the  lead 
nitrate  solution  place  a  piece  of  zinc,  and  allow  it  to 
remain  undisturbed  as  long  as  possible.  The  product  is 
lead. 

QUESTIONS 

1.  What  is  the  effect  of  the  continual  absorption  of  small 
quantities  of  lead  upon  the  body,  and  what  are  the  antidotes  ? 

2.  What  is  the  valence  of  lead  in  PbO,  PbCla,  PbS,  Pb(NO,)s, 
and  PbO9  ? 

PROBLEMS 

114.  If  50  grams  of  litharge  (PbO)  contain  3.5862  grams  of 
oxygen,  what  is  the  atomic  weight  of  lead  ? 

115.  What  are  the  name  and  formula  of  a  compound  of  the 
following  composition  ? 

Lead 77.52  per  cent. 

Carbon 4.49      " 

Oxygen 17.98      " 


CHAPTERS    XXX    AND    XXXI 

No  experiments. 

QUESTIONS 

1.  What  reasons  have  we  for  the  belief  that  red  and  colorless 
phosphorus  are  two  forms  of  the  same  element  ? 

2.  Discuss  the  luminosity  of  colorless  phosphorus.     How  is 
the  luminosity  affected  by  the  composition  of  the  gas  in  contact 
with  the  substance  ? 

PROBLEMS 

116.  How  many  liters  of  oxygen  are  needed  to  bum  93  grams 
of  phosphorus  to  P2OB  ? 

117.  20  grams  of  phosphorus  are   burned  in  a  vessel  from 
which  nothing  is  allowed  to  escape.     How  much  will  the  vessel 
increase  in  weight  ? 

118.  What  volume  of  air  is  needed  to  burn  124  grams  of 


ELEMENTARY  CHEMISTRY 

phosphorus  to  PaO&  ?     Assume  that  air  contains  21  per  cent  by 
volume  of  oxygen. 

119.  If  4  grams  of  phosphorus  when  burned  yield  9.16  grams 
of  P»O6,  what  is  the  atomic  weight  of  phosphorus  ? 


CHAPTERS   XXXII    AND    XXXIII 

ARSENIC  AND   ANTIMONY 

EXPERIMENT  101. — Arsenic. — Examine  the  element 
and  record  its  properties.  Heat  a  fragment  of  it  in  a 
tube  of  hard  glass,  closed  at  one  end.  Does  it  melt  ?  Does 
it  vaporize?  The  two  sublimates  (steel-gray  and  black) 
are  different  allotropic  forms  of  arsenic.  Heat  a  frag- 
ment of  arsenic  the  size  of  the  head  of  a  pin  on  charcoal 
with  the  blowpipe  flame.  The  product  is  arsenious  oxide. 

EXPERIMENT  102. — Arsenious  oxide. — Examine  the 
substance  and  compare  it  with  the  element  arsenic.  Heat 
a  trace  of  arsenious  oxide  in  a  sealed  tube.  (?)  Examine 
the  sublimate  with  a  lens.  What  is  the  shape  of  the  crys- 
tals? 

Dissolve  about  0.1  gram  of  arsenious  oxide  by  boiling 
it  gently  in  a  test-tube  with  dilute  hydrochloric  acid. 
Dilute  the  liquid  to  100  c.c.  Into  a  portion  of  this  liquid 
pass  hydrogen  sulphide  gas — a  slow  succession  of  bubbles. 
Use  a  Kipp  generator  as  a  source  of  the  gas,  or  generate 
it  in  a  wide  test-tube,  as  directed  in  Experiment  110. 
Work  under  the  hood,  and  do  not  inhale  the  hydrogen 
sulphide.  Describe  what  takes  place,  and  write  the  equa- 
tion. 

Place  a  clean  piece  of  sheet  copper  about  2  cm.  square 
in  a  dish,  fill  the  dish  with  water,  add  about  1  c.c.  of 
hydrochloric  acid,  and  heat  almost  to  boiling.  Is  there 
any  action?  Add  1  c.c.  of  your  arsenic  solution  to  the 


ARSENIC  AND  ANTIMONY  113 

liquid,  and  continue  heating  for  ten  minutes.  The  deposit 
on  the  copper  is  arsenic.  Remove  the  copper,  dry  it  care- 
fully with  filter  paper,  roll  it  up,  and  place  it  in  a  tube 
of  hard  glass,  -sealed  at  one  end.  Heat  it  gently.  ( ?) 
Look  for  a  sublimate  with  a  lens. 

Mix  about  0.1  gram  of  arsenious  oxide  with  twice  its 
weight  of  powdered  charcoal,  and  heat  the  mixture  in  a 
tube  of  hard  glass,  sealed  at  one  end.  Introduce  the 
mixture  with  a  paper  trough.  The  upper  part  of  the  tube 
must  be  clean.  Describe  and  explain  the  result. 

EXPERIMENT  103. — Antimony. — Examine  the  element. 
Is  it  a  metal?  How  does  it  behave  under  the  hammer? 
Heat  a  fragment  on  charcoal  with  the  blowpipe.  The 
product  is  antimonious  oxide,  Sb406.  Let  the  globule  of 
melted  antimony  fall  while  still  red  hot  on  a  sheet  of 
paper.  (?) 

QUESTION 

Explain  in  detail  what  is  meant  by  the  statement  that  arsenic 
and  antimony  stand  on  the  border-line  between  metals  and  non- 
metals.  Consider  both  physical  and  chemical  properties  in  your 
answer. 

PROBLEMS 

120.  How  much  charcoal  must  be  used  to  reduce  132  grams  of 
arsenious   oxide  ?      What  volume  of   carbon  monoxide  will  be 
given  off  ? 

As4O6  +  6C  =  6CO  +  As4. 

121.  In  a  case  of  poisoning  the  arsenic  from  the  body  was 
converted   into   arsenious   sulphide,    As28s,    and   weighed.      The 
weight  of  the  As,St  was  0.82  gram.     What  quantity  of  arsenious 
oxide  had  been  administered  ? 

122.  What  is  the  weight  of  5.6  liters  of  arsine  at  0°  and 
760  mm.  ? 

123.  Calculate  the  formula  of  a  compound  of  the  following 
composition  : 

Arsenic 48.39  per  cent. 

Sulphur 51.61      " 


114  ELEMENTARY  CHEMISTRY 

CHAPTEE   XXXIV 

CHROMIUM 

EXPERIMENT  104. — Chrome  alum.  Preparation  of  po- 
tassium chromate,  K2Cr04. — Examine  chrome  alum.  What 
is  the  form  of  the  crystals?  Powder  a  little.  With  a 
portion  of  the  powder  try  the  solubility  in  water.  What 
is  the  color  of  the  liquid  ?  Chrome  alum  contains  chromi- 
um sulphate,  Cr2(S04)3,  and  in  it  and  the  other  chromium 
salts  chromium  plays  the  chemical  role  of  a  trivalent 
metal,  like  aluminium. 

Mix  the  rest  of  the  powdered  chrome  alum  with  po- 
tassium nitrate  and  potassium  carbonate  in  a  clean  mor- 
tar. The  mixture  should  contain  equal  quantities  of  the 
three  substances.  Fuse  the  mixture  on  a  piece  of  platinum 
foil,  held  in  forceps.  Notice  the  color  of  the  melted  sub- 
stance. It  now  contains  potassium  chromate,  K2Cr04,  in 
which  the  chromium  plays  the  role  of  a  non-metal,  like 
sulphur. 

Dissolve  the  melted  substance  in  water  and  carefully 
add  acetic  acid  until  litmus  paper  is  turned  red.  Notice 

the  color  of  the  liquid.     It  is  due  to  the  negative  ion 

—  —  +++ 

Cr04.     What  color  did  positive  ions  Or  give  to  water 

when  you  dissolved  chrome  alum  in  it?  Divide  the  liquid 
into  two  portions.  To  half  of  it  add  solution  of  lead 
nitrate.  Complete  the  equation — 

K2Cr04  +  Pb(N03)2  =  PbCr04  + 

Lead  chromate 

To  the  rest  of  the  liquid  add  either  nitric  or  sulphuric 
acid,  drop  by  drop,  until  the  yellow  changes  to  red.  The 
red  liquid  contains  potassium  dichromate,  K2Cr207.  Com- 
plete the  equation — 

2K2Cr04  +  2HN03  =  K2Cr207  +~  ...  +  ... 


SULPHUR  115 

EXPERIMENT  105. — Preparation  of  potassium  chromate 
from  potassium  dichromate. — Dissolve  5  grams  of  potas- 
sium dichromatc  in  water  in  a  porcelain  dish.  Add  a 
solution  of  potassium  hydroxide  until  the  liquid  is  yellow. 
Evaporate  to  small  volume  and  let  the  solution  crystallize. 
The  product  is  potassium  chromate.  Complete  the  equa- 
tion— 

K2Cr207  +  2KOH  =...  +  ... 

•  EXPERIMENT  106. — Preparation  of  chromic  chloride 
from  potassium  chromate. — Powder  a  little  potassium 
chromate,  and  heat  it  gently  in  a  test-tube  under  the  hood 
with  strong  hydrochloric  acid.  What  gas  escapes?  What 
is  the  change  in  the  color  of  the  liquid  ?  It  now  contains 
chromic  chloride,  CrCl3,  in  which  chromium  acts  like  a 
metal,  chemically.  Supply  the  lacking  numbers  in  the 
equation — 

K2Cr04  +  8HC1  =  CrCl3  +  KC1  +  01  +  H20 

QUESTIONS 

1.  Briefly  state  the  steps  by  which  you  have  transformed  a 
chromic  salt  into  a  chromate.     A  chromate  into  a  chromic  salt. 

2.  How  can  potassium  chromate  be  converted  into  potassium 
dichromate  ?     Potassium  dichromate  into  potassium  chromate  ? 

PROBLEM 

124.  Calculate  the  percentage  composition  of  clvromite, 
CraOaFeO. 

CHAPTER   XXXV 

SULPHUR 

EXPERIMENT  107. — Different  forms  of  sulphur. — Ex- 
amine roll  sulphur.     Is  it  brittle  or  malleable?  dense  or 
light?    Rub  a  piece  on  the  coat-sleeve  and  bring  it  near 
30 


116  ELEMENTARY  CHEMISTRY 

small  pieces  of  paper.  Does  it  conduct  electricity?  Hold 
a  piece  firmly  in  the  hand  close  to  the  ear  and  notice  the 
crackling  noise  caused  by  the  portions  of  sulphur  next  the 
hand,  expanding  as  they  become  warm  and  cracking  away 
from  the  other  portions,  which  remain  cold.  This  shows 
that  it  is  a  bad  conductor  of  heat. 

Examine  flowers  of  sulphur,  and  record  its  properties. 
Powder  about  2  grams  of  roll  sulphur  and  shake  it  up 
with  3-5  c.c.  carbon  disulphide  in  a  dry  test-tube.  Cau- 
tion: Keep  the  carbon  disulphide  bottle  corked,  and  avoid 
the  vicinity  of  flame.  When  the  sulphur  has  dissolved, 
pour  the  liquid  into  a  dish  and  let  it  evaporate  spon- 
taneously. Examine  the  crystals  with  a  lens.  They  con- 
sist of  a-sulphur. 

Fill  the  smallest  size  Hessian  crucible  with  crushed 
roll  sulphur,  place  it  in  a  ring  of  your  stand,  and  apply 
heat  until  the  sulphur  has  melted.  Allow  to  cool.  Just 
when  the  surface  begins  to  solidify  take  the  crucible  in 
forceps  and  pour  out  the  liquid  portion  into  a  pan  of 
water.  Let  it  cool,  and  examine  the  interior.  It  is  cov- 
ered with  crystals  of  ^-sulphur.  Make  a  drawing  of  a 
mass  of  the  crystals.  Record  their  properties.  Notice 
that  they  are  different  in  color  as  well  as  shape  from  the 
crystals  of  a-sulphur.  Preserve  some  crystals  overnight 
and  explain  the  change  which  takes  place. 

When  roll  sulphur  is  first  made,  does  it  consist  of 
a-sulphur  or  ^-sulphur?  Which  does  it  consist  of  after 
being  preserved  for  a  time? 

In  what  respects  does  the  change  of  a-sulphur  into 
/9-sulphur  by  heat  (p.  236,  Part  I)  resemble,  and  in  what 
respects  differ,  from  the  change  of  ice  into  water? 

EXPERIMENT  108. — Fusion  of  sulphur.  Soft  sulphur. 
— Fill  a  wide  test-tube  about  one-quarter  with  crushed 
roll  sulphur,  hold  it  with  a  paper  holder,  and  heat  slowly. 
Obtain  the  three  states  of  fusion  described  on  page  235, 


SULPHUR  117 

Part  I.  The  three  conditions  should  be  perfectly  distinct. 
Let  the  sulphur  in  the  second  liquid  stage  cool  slowly, 
and  note  that  it  passes  through  the  changes  in  the  re- 
verse order. 

Now  heat  the  sulphur  until  it  boils  vigorously,  and  pour 
the  liquid  in  a  thin  stream  into  a  bottle  of  cold  water, 
lighting  the  sulphur  vapor  at  the  mouth  of  the  tube  while 
pouring.  Examine  the  product  and  describe  it.  Preserve 
some  of  it.  (?)  Is  it  a  stable  condition  of  sulphur? 

EXPERIMENT -109. — Union  of  sulphur  with  metals. — 
Fill  the  same  test-tube  one-quarter  with  sulphur,  and  heat 
to  boiling.  Into  the  tube  throw  a  little  powdered  iron.  (?) 
Straighten  a  thin  iron  wire,  heat  one  end  of  it  to  red- 
ness, and  quickly  place  it  in  the  sulphur  vapor.  Cut  a 
thin  strip  of  copper  and  hold  it  in  the  tube. 

How  does  sulphur  behave  with  metals?  What  is  a 
sulphide  ? 

EXPERIMENT  110. — Hydrogen  sulphide. — Use  the  hood. 
Hydrogen  sulphide  is  poisonous,  and  must  not  be  inhaled. 
Cover  the  bottom  of  a  gas-generating  bottle  like  the  one 
used  in  making  hydrogen  with  lumps  of  iron  sulphide, 
FeS.  Incline  the  bottle,  and  slide  the  solid  in  slowly. 
Cover  with  water,  and  add  hydrochloric  acid  through  the 
funnel-tube.  Collect  several  wide  test-tubes  full  of  the 
gas  over  water  as  warm  as  you  can  work  with.  (Why  is 
it  necessary  to  use  warm  water?)  Cork  one  of  the  test- 
tubes,  transfer  it  to  a  pan  of  cold  water,  remove  the  cork, 
and  shake  the  tube,  keeping  its  mouth  under  water.  Is 
the  gas  soluble  in  water?  With  another  test-tube  try  its 
combustibility.  Meanwhile,  allow  the  hydrogen  sulphide 
to  bubble  through  a  bottle  or  beaker  containing  at  least 
\  liter  of  cold  water.  Hold  the  exit  tube  of  your  appa- 
ratus near  a  silver  coin.  A  clean  piece  of  sheet  copper. 
A  paper  wet  with  a  solution  of  lead  acetate.  Record  and 
explain  the  results. 


118  ELEMENTARY  CHEMISTRY 

Put  the  exit  tube  back  in  the  water  and  continue  pass- 
ing the  gas  through  for  ten  minutes.  Meanwhile,  prepare 
very  dilute  solutions  of  copper  sulphate,  lead  nitrate,  cad- 
mium chloride,  tartar  emetic  (an  antimony  compound), 
and  sodium  chloride.  Stop  the  production  of  the  gas 
by  filling  the  bottle  with  water  and  pouring  away  the 
liquid.  Has  the  water  in  the  beaker  the  odor  of  the  gas? 
Use  it  to  study  the  effect  of  hydrogen  sulphide  upon  the 
solutions  you  have  prepared.  Complete  the  equations: 


Fb(N08) 


Does  the  hydrogen  sulphide  affect  the  sodium  Chloride 
solution?  Other  metals  —  e.  g.,  potassium,  calcium,  and 
magnesium  —  behave  like  sodium.  How  could  you  sepa- 
rate one  of  these  metals  from  lead,  cadmium,  or  copper? 

QUESTIONS 

1.  What  is  the  natural  state  of  sulphur  at  ordinary  tempera- 
tures ?    At  100°  ?     Give  facts  to  support  your  answer. 

2.  Why  is  hydrogen  sulphide   so  much  used  in  analytical 
chemistry  ? 

PROBLEMS 

125.  2  grams  of  crystallized  copper  sulphate  (CuSO45HaO)  are 
dissolved  in  water,  and  it  is  required  to  precipitate  all  the  copper 
as  CuS.     How  much  iron  sulphide,  and  how  much  hydrochloric 
acid  containing  25  per  cent  HC1,  are  needed  to  generate  enough 
HaS  for  the  purpose  ? 

126.  (a)  What  volume  of  hydrogen  sulphide  is  produced  when 
17.6  grams  of  FeS  are  dissolved  in  HC1  ?     (ft)  From  what  weight 
of  mercuric  chloride  dissolved  in  water  will  this  quantity  of  H8S 
precipitate  the  mercury  as  mercuric  sulphide  ? 

HgCla  +  HaS  =  HgS  +  2HC1. 

127.  What  volume  of  air  is  needed  to  burn  500  grams  of  sul- 
phur to  SOa  ?    Assume  that  air  contains  21  per  cent  of  oxygen  by 
volume. 


SULPHUR  119 

128.  What  volume  of  H2S  escapes  when  5  grams  of  iron  sul- 
phide are  dissolved  in  HC1  ? 

129.  How  much  iron  sulphide  is  needed  to  make  50  liters  of 
HaS? 

In  working  with  sulphur  dioxide,  use  the  hood.  The 
gas  is  irritating,  and  must  not  be  inhaled. 

EXPERIMENT  111. — Sulphur  dioxide,  S02. — Fill  a  small 
flask  one-third  with  a  strong  solution  of  sodium  acid  sul- 
phite, NaHS03,  and  allow  strong  sulphuric  acid  to  drop 
slowly  into  it  from  a  separating  funnel  which  passes 
through  one  hole  in  the  cork.  Lead  away  the  S02  by 
a  tube  passing  through  the  other  hole,  and  collect  it  by 
downward  displacement  in  dry,  covered  vessels.  Collect 
two  wide  test-tubes  and  two  bottles  of  the  gas.  Cover  the 
test-tubes  with  paper  and  the  bottles  .with  glass  plates. 

Investigate  its  solubility  by  placing  a  test-tube  filled 
with  it  mouth  down  in  water.  Kemove  the  test-tube  from 
the  pan  by  slipping  a  glass  plate  under  it,  and  add  a  few 
drops  of  litmus  solution  to  the  liquid  in  the  tube.  The 
reddening  of  the  litmus  shows  the  presence  of  sulphurous 
acid.  How  has  it  been  formed? 

Place  a  few  drops  of  litmus  solution  in  a  test-tube 
filled  with  the  gas,  cork  the  tube,  shake  it,  and  let  stand. 
What  is  the  first  action  of  the  gas  on  the  litmus?  Why? 
What  is  the  final  action?  Sulphur  dioxide  is  much  used 
for  bleaching. 

Lower  a  lighted  candle  into  a  bottle  of  the  gas.  (?) 
Place  some  lead  dioxide  in  a  clean  deflagrating  spoon  and 
warm  it  gently.  Lower  it  into  a  bottle  of  sulphur  di- 
oxide. The  product  is  lead  sulphate — 

Pb02  +  S02  =  PbS04 
Examine  it,  and  compare  it  with  lead  dioxide. 

Pass  sulphur  dioxide  from  your  generator  through 
100  c.c.  of  water,  until  you  have  a  strong  solution.  Test 
the  liquid  with  litmus  paper  (red  and  blue).  Add  a 


120  ELEMENTARY  CHEMISTRY 

dilute  solution  of  sodium  hydroxide  to  it  drop  by  drop, 
stirring  constantly  until  litmus  paper  is  just  turned  blue. 
Then  evaporate  to  dryness  in  a  clean  dish.  The  residue 
is  sodium  sulphite,  Na2S03,  formed  thus — 

SNaOH  +  H2S03  =  Na2S03  +  2H20 

Examine  it.  How  does  it  behave  with  hydrochloric 
acid?  with  sulphuric  acid?  What  gas  is  produced  when 
the  acids  act  upon  it?  Take  some  sodium  sulphite  from 
the  bottle  and  see  if  it  behaves  in  the  same  way  with 
acids. 

Most  sulphites  are  insoluble  in  water.  Mix  water  so- 
lutions of  sodium  sulphite  and  calcium  chloride.  The 
visible  product  is  calcium  sulphite,  CaS03.  Write  the 
equation.  Is  it  soluble  in  water?  in  hydrochloric  acid? 

If  time  permits,  pass  a  slow  current  of  sulphur  diox- 
ide from  your  generator  through  a  U-shaped  tube,  which 
is  embedded  in  a  mixture  of  equal  parts  of  ice  and  salt. 
Liquid  sulphur  dioxide  collects.  Pour  some  of  it  on  a 
little  water  in  a  porcelain  dish.  Use  the  hood.  The  water 
will  be  frozen  by  the  rapid  evaporation  of  the  sulphur 
dioxide. 

EXPERIMENT  112. — Sulphuric  acid, — Examine  sul- 
phuric acid.  How  does  it  compare  with  water  as  re- 
gards density  and  consistency?  Slowly  pour  sulphuric 
acid  in  a  thin  stream,  stirring  constantly  into  5  c.c.  of 
water  in  a  beaker  until  you  have  added  an  equal  volume 
of  the  acid.  Care.  Never  add  water  to  sulphuric  acid. 
Always  add  the  acid  slowly  to  the  water,  stirring  con- 
stantly. Notice  the  production  of  heat.  Make  some  dilute 
sulphuric  acid,  and  write  a  word  with  it  on  a  sheet  of 
paper,  using  a  glass  rod  as  a  pen.  Heat  the  paper  over 
the  flame,  taking  care  not  to  set  fire  to  it.  Place  a  splinter 
or  a  match-stick  in  a  little  strong  sulphuric  acid  in  a 
test-tube. 


SULPHUR  121 

Make  a  little  strong  hot  solution  of  sugar  and  place 
not  more  than  3  c.c.  of  it  in  a  test-tube.  Slowly  add 
strong  sulphuric  acid  until  a  decided  result  is  obtained. 

Wood  and  paper  are  chiefly  composed  of  a  substance 
called  cellulose,  whose  formula  is  C6H1005.  The  formula 
of  sugar  is  C12H22011.  Bearing  in  mind  the  strong  at- 
traction of  sulphuric  acid  for  water,  explain  the  three 
results  just  obtained. 

Eepeat  the  reaction  of  barium  chloride  with  sulphuric 
acid.  (Experiment  81.) 

Look  through  your  notes  from  the  beginning  for  all 
cases  in  which  sulphuric  acid  has  been  used,  and  make 
a  table  in  your  notes  containing  the  equations  for  the 
various  reactions  in  which  you  have  employed  it.  Ar- 
range the  table  thus: 

Reactions  of  Sulphuric  Acid  with  Various  Substances 


Zn  H 

h  HaS04 

—  ZnS04  4 

Ha 

Sodium  nitrate              . 

2NaN03  H 

h  H,S04 

-  NaaS04 

•f  2HNO, 

QUESTIONS 

1.  What  is  the  most  recent  process  for  the  production  of  sul- 
phuric acid  ? 

2.  What  is  the  lead-chamber  process  for  the  production  of 
sulphuric  acid  ? 

PBOBLEMS 

130.  How  much  sulphuric  acid  can  be  made  from  4  tons  of 
sulphur  ? 

131.  How  much  sulphuric  acid  can  be  made  from  40  tons  of 
pyrite,  FeS9  ? 

132.  The  density  of  sulphuric  acid  is  1.84.     How  much  sul- 
phur is  there  in  100  c.c.  of  it  ? 

133.  How  many  tons  of  sulphuric  acid  can  be  made  from  100 
tons  of  pyrite  containing  48  per  cent  of  sulphur  ? 


122  ELEMENTARY  CHEMISTRY 

134.  1.8752  grams  of  cobalt,  when  converted  into  cobalt  sul- 
phate, yielded  4.9472  grams.  What  is  the  atomic  weight  of 
cobalt  ?  Assume  S  =  32,  O  =  16.  The  formula  of  cobalt  sulphate 
is  CoSO4. 


CHAPTEE   XXXVI 
No  experiments. 

CHAPTER   XXXVII 

FLUORINE 

EXPERIMENT  113. — Hydrofluoric  acid. — Mix  in  a  lead 
dish — under  the  hood — some  powdered  calcium  fluoride 
with  about  an  equal  weight  of  strong  sulphuric  acid.  Sup- 
port the  dish  in  a  ring  or  on  a  pipe-stem  triangle,  and 
apply  a  gentle  heat.  The  flame  should  not  touch  the  dish. 
Do  not  inhale  the  gas  given  off.  It  is  poisonous.  Hold 
over  the  dish  a  piece  of  red  and  a  piece  of  blue  litmus 
paper.  (?)  Wet  a  glass  rod  with  ammonia.  Be  careful 
not  to  let  any  ammonia  fall  into  the  dish,  since  this  would 
cause  an  explosion.  Remembering  the  behavior  of  am- 
monia  with  HC1,  explain  its  behavior  with  HF.  /Cover  the 
dish  with  a  clean  plate  of  window-glass,  arid  let  it  stand 
some  hours,  if  possible.  Remove  the  glass,  wash  it  with 
water,  and  examine.  Clean  the  dish  by  scraping  the  con- 
tents into  a  jar  with  an  old  k»i£e~er-&le.  ^ [^ U  U 

Warm  a  glass  plate  by  placing  it  on  a  ring  of  your 
stand  gnd  moving  about  a  small  flame  at  least  5  cm. 
below  the  plate.  If  the  flame  is  brought  nearer  the  plate 
it  will  crack  it.  Place  some  paraffin  on  the  warm  plate, 
and  allow  the  melted  paraffin  to  cover  it  completely.  Let 
the  plate  cool  and  scratch  a  word  or  some  lines  through 
the  wax  with  a  pin  or  a  knife-blade.  Take  the  plate  to 


Jb-^ifoji/t  Ws, 


BROMINE  AND  IODINE 

PL  o^v -  & 

the  hood,  cover  the  word  with  a  piece  of  filter  paper,  and 
carefully  wet  the  paper  with  hydrofluoric  acid.  Be  care- 
ful. Hydrofluoric  acid  produces  poisoned  wounds  on  the 
skin,  and  its  vapor  is  poisonous.  Place  a  paper  containing 
a  plainly  written  caution  by  the  side  of  the  plate,  so  that 
no  one  may  touch  it  and  be  injured,  and  let  it  stand  under 
the  hood  for  as  long  a  time  as  possible.  Then  take  away 
the  paper  with  forceps,  and  throw  it  into  the  waste  jar. 
Take  the  plate  with  forceps  and  place  it  in  a  pan  of  boil- 
ing water.  This  will  remove  both  the  hydrofluoric  acid 
and  the  wax.  Wipe  it  with  a  towel  and  examine.  Write 
the  equations  between  calcium  fluoride  and  sulphuric  acid 
and  between  silica  and  hydrofluoric  acid. 

PROBLEMS 

135.  How  much  calcium  fluoride  and  how  much  sulphuric 
acid  containing  96  per  cent  HaSO4  are  needed  to  make  12  grams 
of  pure  hydrofluoric  acid  ? 

136.  How  much  calcium  sulphate  and  how  much  hydrofluoric 
acid  are  formed  when  50  grams  of  calcium  fluoride  are  heated  with 
sulphuric  acid  ? 


CHAPTER   XXXVIII 

BROMINE  AND  IODINE 

EXPERIMENT  114. — Bromine. — Use  the  hood  in  work- 
ing with  bromine.  The  vapor  is  extremely  irritating. 
The  liquid  must  not  be  gotten  upon  the  flesh  or  clothing. 
Powder  1  gram  of  potassium  bromide  and  mix  it  with  2 
grams  of  powdered  manganese  dioxide.  Place  the  mix- 
ture in  a  wide  test-tube,  and  fill  the  tube  one-quarter  with 
cold  dilute  sulphuric  acid,  1  volume  acid  to  4  volumes 
water.  Clamp  the  tube  in  an  inclined  position,  connect 
it  with  a  well-fitting  cork  and  delivery  tube,  dipping  into 


124  ELEMENTARY  CHEMISTRY 

an  empty  test-tube  surrounded  by  cold  water.  Heat  gen- 
tly. The  product  is  bromine.  Examine  it.  The  equation 
is  similar  to  that  given  for  iodine,  on  p.  258,  Part  I. 
Write  it. 

JExamine  bromine,  attending  especially  to  color,  den- 
sity, and  volatility.  Place  three  dry  test-tubes  in  a  rack 
under  the  hoo'd.  Put  a  little  bromine  in  each  (1  cm. 
deep),  ^otice  the  color  of  the  vapor.  Into  one  test-tube 
drop  a  piece  of  tin,  into  another  a  piece  of  antimony. 
Use  pieces  a  little  larger  than  the  head  of  a  pin,  and  be 
careful.  Is  bromine  an  energetic  element.  (Nearly  fill 
the  third  test-tube  with  water  and  cork  it  tightly.  Shake 
it.  Notice  the  color  of  the  product,  which  is  called  bro- 
mine-water. Does  bromine-water  bleach?  Try  it  on  a 
little  water  faintly  colored  with  litmus. 

Dissolve  a  little  potassium  bromide  in  water  in  a  wide 
test-tube.  Add  a  layer  of  chloroform  2  cm.  deep,  cork 
the  tube,  and  shake  it.  Do  bromine  ions  give  any  color 
to  water  or  chloroform?  Now  add  chlorine-water  to  the 
liquid  until  the  tube  is  nearly  full,  cork  it,  and  shake  it 
for  five  minutes.  The  color  of  the  chloroform  is  due  to 
dissolved  bromine,  Br2,  not  bromine  ions.  Write  the  equa- 
tion, first  in  the  ordinary  way;  second,  using  the  idea  of 
ions.  The  chloroform  and  the  water  take  no  part  in  the 
chemical  change,  and  must  not  appear  in  the  equation. 

Pour  away  the  water  as  far  as  possible  without  losing 
any  of  the  chloroform,  fill  up  the  tube  with  water,  shake 
and  pour  off  again  to  get  rid  of  all  chlorine.  A  solution 
of  Br2  in  chloroform  remains  in  the  tube.  Add  more 
water  and  a  crystal  of  potassium  iodide,  KI,  and  shake 
again.  Notice  the  change  in  the  color  of  the  chloroform. 
It  now  contains,  instead  of  Br2,  I2  dissolved.  Write  the 
equation,  both  from  the  ordinary  standpoint  and  from  the 
standpoint  of  ions. 

Discuss  this  experiment  at  length  in  your  notes,  and 


BROMINE  AND  IODINE  125 

show  from  it  that  the  tendency  of  these  three  elements  to 
exist  as  ions  decreases  in  the  order  Cl,  Br,  I.  Where 
would  F  come  in  this  series? 

EXPERIMENT  115. — Hydrobromic  acid. — Place  a  few 
crystals  of  potassium  bromide  in  a  wide  test-tube  and 
cover  with  a  mixture  of  3  volumes  of  sulphuric  acid  and 
1  volume  of  water.  Clamp  the  tube  in  an  inclined  posi- 
tion, insert  a  cork  with  a  delivery  tube,  and  collect  the 
gas  by  downward  displacement  in  dry  wide  test-tubes. 
Use  a  gentle  heat.  Try  the  solubility  in  water.  Try  the 
action  upon  red  and  blue  litmus  paper.  Let  fall  a  drop 
of  ammonia  into  a  tube  of  the  gas.  Make  a  comparison 
between  HBr  and  HC1. 

EXPERIMENT  116. — Iodine. — Place  in  a  mortar  1  gram 
of  potassium  iodide  and  2  grams  of  manganese  dioxide. 
Grind  thoroughly  and  transfer  the  mixture  to  a  porcelain 
dish.  Add  enough  of  a  cold  mixture  of  1  volume  strong 
sulphuric  acid  and  2  volumes  water  to  cover  the  powder. 
Cover  the  dish  with  a  funnel  and  apply  a  gentle  heat. 
When  the  funnel  is  well  coated  with  iodine-crystals,  stop 
the  experiment  and  scrape  off  the  crystals.  Use  them  for 
the  following  experiments,  or,  if  the  quantity  is  insuffi- 
cient, obtain  a  little  iodine  from  the  bottle. 

Examine  iodine  with  respect  to  color,  luster,  odor,  and 
density.  Warm  a  dry  wide  test-tube  gently,  and  drop  a 
crystal  of  iodine  into  it.  Invert  the  tube.  What  are  the 
properties  of  iodine  vapor?  What  about  its  density?  If 
the  formula  of  iodine  is  I2,  what  must  be  the  density  of 
its  vapor  referred  to  hydrogen?  referred  to  air? 

Use  the  same  test-tube  to  test  the  solubility  of  iodine 
in  water.  ( ?)  Use  the  tip  of  a  knife-blade  filled  with 
powdered  iodine — not  more.  Add  two  crystals  of  potas- 
sium iodide  and  shake.  (  ?)  Preserve  the  solution. 

Grind  about  0.5  gram  of  starch  with  a  little  water  in 
a  mortar.  Slowly  pour  the  liquid  into  100  c.c.  of  boiling 


126  ELEMENTARY  CHEMISTRY 

water  in  a  clean  dish.  The  liquid  must  be  colorless.  It 
remains  colorless  when  a  few  drops  of  it  are  mixed  with  a 
little  potassium-iodide  solution  in  a  test-tube.  Try  it. 
Now  half  fill  a  wide  test-tube  with  water,  add  a  little  of 
your  starch  solution,  and  a  drop  of  your  solution  of  iodine. 
Notice  the  color,  which  can  be  used  as  a  test  both  for 
iodine  and  starch.  The  color  disappears  when  the  tube  is 
heated,  but  reappears  on  cooling.  Try  this. 

Place  about  3  c.c.  of  chloroform  in  a  test-tube,  half 
fill  the  test-tube  with  water,  and  add  a  drop  or  two  of 
your  solution  of  iodine  in  potassium  iodide.  Shake  the 
tube.  The  color  of  the  chloroform  solution  can  be  used 
as  a  test  for  iodine. 

Carefully  examine  potassium  iodide.  What  are  the 
corresponding  bromine  and  chlorine  compounds?  Does 
it  resemble  them?  Add  a  few  drops  of  potassium  iodide 
solution  to  a  little  solution  of  silver  nitrate.  The  visible 
product  is  silver  iodide,  Agl.  Write  the  equation.  Make 
some  silver  bromide  in  the  same  way.  Do  they  resemble 
silver  chloride?  If  you  can  not  recall  the  properties  of 
silver  chloride,  make  a  little  by  mixing  solutions  of  silver 
nitrate  and  sodium  chloride.  Use  half  of  each  product  to 
study  the  solubility  of  the  three  substances  in  ammonia. 
Stand  the  other  half  in  direct  sunlight,  or  the  brightest 
light  accessible.  Eecord  and  discuss  the  results. 

PROBLEMS 

137.  How  much  potassium  bromide  is  required  to  make  50 
grams  of  bromine  ? 

138.  1.8  grams  of  potassium  bromide  are  heated  with  dilute 
sulphuric  acid.     What  volume  of  hydrobromic-acid  gas  at  26°  is 
liberated  ? 

139.  What  volume  of  hydrobromic-acid  gas  can  be  made  from 
160  c.c.  of  liquid  bromine  according  to  the  method  described  on 
p.  257,  Part  I  ?    Assume   that   bromine   is  3  times  as  dense  as 
water. 


IRON  127 

140.  Manganese  dioxide  is  heated  with  hydrochloric  acid,  and 
the  chlorine  passed  in  a  solution  of  potassium  iodide.     How  much 
iodine  will  be  set  free  by  the  chlorine  evolved  when  12  grams  of 
manganese  dioxide  are  used  ? 

141.  Under  the  same  circumstances  as  in  Problem  138,  how 
much  manganese   dioxide   is  needed  to  liberate  63.5  grams  of 
iodine  ? 

142.  Under  the  same  conditions  as  in  the  two  preceding 
problems,  how  much  iodine  will  be  set  free  when  43.5  grams  of 
manganese  dioxide  are  used  ?  , 


CHAPTER   XXXIX 

IRON 

>^  EXPERIMENT  117. — Reduction  of  ferric  oxide  by  hy- 
arogen. — Place  a  little  ferric  oxide  in  a  bulb-tube  and  pass 
hydrogen,  dried  by  passing  through  a  U-tube  filled  with 
lumps  of  calcium  chloride,  over  it.  It  is  best  to  use  a 
Kipp  generator.  If  you  employ  an  ordinary  gas-bottle, 
it  is  necessary  to  wait  until  all  air  is  expelled,  and  to 
wrap  the  gas-bottle  and  cork  in  a  towel  before  applying 
heat  to  the  bulb.  Heat  the  bulb,  gently  at  first.  What 
are  the  two  products?  Let  the  bulb  cool  completely, 
shake  out  the  solid  material  on  a  paper,  and  examine  it. 
Is  it  attracted  by  the  magnet?  Is  it  combustible? 

EXPERIMENT  118. — Spontaneous  oxidation  of  finely 
divided  iron. — Heat  in  a  sealed  tube  of  hard  glass  some 
ferrous  oxalate  to  redness,  keeping  the  finger  loosely  over 
the  open  end  of  the  tube  to  prevent  the  entrance  of 
air.  Cork  the  tube  tightly,  let  it  cool,  and  shake  out 
the  black  iron-powder  into  a  plate  or  a  dry  dish.  ( ?) 
Clearly,  finely  divided  iron  oxidizes  more  readily  than 
compact  iron.  You  have  noticed  frequently  in  your  work 
that  powdering  a  substance  makes  it  dissolve  more  rap- 


128  ELEMENTARY  CHEMISTRY 

idly.  The  reason  is  the  same  in  the  two  cases.  What 
is  it? 

Examine  the  most  important  ores  of  iron,  hematite 
(Fe203),  magnetite  (Fe304),  and  limonite,  which  contains 
both  ferric  oxide  and  ferric  hydroxide.  Also  examine  py- 
rite  (FeS2),  which  is  important  on  account  of  its  sul- 
phur. 

EXPERIMENT  119. — Ferric  salts. — Dissolve  a  little  fer- 
ric chloride,  FeCl3,  in  50  c.c.  of  water  in  a  beaker.  Use 
small  portions  to  investigate  the  action  of  the  following 
substances  upon  it: 

1.  Ammonia  water,  NH4OH,  precipitates  ferric  hy- 
droxide, Fe(OH)3.     Write  the  equation  and  record  the 
properties  of  the  substance. 

2.  Potassium  tliiocyanate,  KCNS  (often  called  potas- 
sium sulphocyanide).     Dissolve  a  little  in  water  for  the 
test.     Notice  that  a  red  coloration  of  ferric  thiocyanate, 
Fe(CNS)3,  is  obtained,  not  a  precipitate. 

3.  Potassium   ferricyanide    (not   ferro  cyanide).     The 

solution  of  this  substance  must  be  made  up 
fresh  and,  before  dissolving,  the  outer  portion 
of  the  crystal  used  should  be  removed  by  run- 
ning water  over  it.  Notice  that  there  is  no 
precipitate,  but  a  dark  coloration. 

EXPERIMENT  120. — Ferrous  salts. — Examine 
some  iron  wire,  and  if  the  wire  contains  any 
rust,   remove   it   by   scraping   with    sandpaper. 
Place  about  0.5  gram  of  the  wire  in  a  small 
flask,  add  100  c.c.  water  and  15  c.c.  hydrochloric 
acid,  and  close  the  flask  with  a  cork  bearing  a 
valve  made  as  shown  in  Fig.  25.     A  glass  tube 
fits  tightly  in  the  cork,  and  is  connected  with  a 
piece  of  rubber  tubing  4  cm.  long,  the  end  of  which  is 
closed  by  a  bit  of  glass  rod.    In  the  middle  of  the  rubber 
tube  is  a  longitudinal  slit  1  cm.  long,  made  with  the  point 


IRON  129 

of  a  sharp  knife-blade.  This  arrangement  permits  the 
escape  of  the  hydrogen  from  the  flask,  but  obstructs  the 
entrance  of  air,  which  would  convert  the  ferrous  chloride 
into  ferric  compounds. 

When  the  iron  is  dissolved,  cool  the  solution  of  fer- 
rous chloride  by  running  water  over  the  flask,  and  investi- 
gate the  behavior  of  small  portions  with  the  same  three 
substances  employed  in  the  preceding  experiment.  Keep 
the  flask  corked. 

Notice  that  ferrous  hydroxide,  obtained  by  the  action 
of  ammonia,  is  at  first  white,  but  rapidly  turns  green,  and 
finally  becomes  rust-colored  ferric  hydroxide.  This  change 
is  due  to  absorption  of  oxygen  from  the  air. 

Potassium  thiocyanate  should  produce  no  color.  If  a 
faint  pink  is  produced,  suggest  a  reason  for  it. 

Potassium  ferricyanide  produces  a  deep  blue  precipi- 
tate, quite  different  in  character  from  the  result  of  its 
action  upon  the  ferric  solution. 

Now  take  a  little  of  your  ferrous  solution  and  add 
bromine-water,  drop  by  drop,  until  the  liquid  is  slightly 
reddish.  Boil  the  liquid.  Notice  the  color  change.  Does 
the  new  color  resemble  that  of  ferric  chloride — the  color 
of  the  ferric  ion  ?  Test  it  with  the  same  three  substances, 
using  separate  small  portions,  and  see  whether  it  is  now 
ferrous  or  ferric.  Write  the  equation  for  the  action  of 
chlorine  on  ferrous  chloride. 

Place  a  little  of  your  ferrous  chloride  solution  in  a 
test-tube,  add  a  few  drops  of  nitric  acid,  and  boil.  Does 
it  seem  to  change  to  ferric  chloride?  Apply  the  three 
tests  and  ascertain. 

Place  a  little  of  your  ferrous  chloride  solution  in  a 
beaker  on  wire  gauze,  and  heat  it  gently  for  15  minutes, 
allowing  free  air  access.  Test  it  for  ferric  compounds 
with  potassium-thiocyanate  solution.  What  is  the  action 
of  air  on  ferrous  salts? 


ISO  ELEMENTARY  CHEMISTRY 

EXPERIMENT  121. — Reduction  of  ferric  to  ferrous  solu- 
tions by  nascent  hydrogen. — Dissolve  not  more  than  0.1 
gram  of  ferric  chloride  in  100  c.c.  of  water.  Carefully 
clean  the  flask  in  which  }^ou  dissolved  the  iron,  place  the 
ferric  chloride  solution  in  it,  incline  the  flask,  and  slide 
in  3  grams  of  granulated  zinc  or  cuttings  of  sheet  zinc, 
add  10  c.c.  of  hydrochloric  acid,  insert  the  valve,  place  the 
flask  on  a  piece  of  asbestos  board,  and  apply  a  gentle  heat. 
The  flask  should  be  heated  at  least  ten  minutes  after  the 
liquid  has  become  colorless.  Determine  whether  the  solu- 
tion is  ferrous  or  ferric  by  the  same  three  tests.  What  is 
the  action  of  nascent  hydrogen  on  ferric  solutions  ?  Write 
the  equation.  Hydrogen  passed  into  a  ferric  solution 
through  a  glass  tube  has  no  effect.  What  is  the  theoretical 
explanation  of  this  fact? 

PBOBLEMS 

143.  What  volume  of  hydrogen  at  13°  and  780  mm.  is  needed 
to  convert  31.5  grams  of  FeaO3  into  iron  ? 

144.  What  volume  of  hydrogen  at  14°  and  740  mm.  is  re- 
quired to  change  20  grams  of  ferric  oxide  into  iron  ? 

145.  If  ferric  oxide,  FeaOs,  contains  70  per  cent  iron  and  30 
per  cent  oxygen,  what  is  the  atomic  weight  of  iron  ? 

146.  If  ferrous  oxide,  FeO,  contains  77.8  per  cent  iron  and 
22.8  per  cent  oxygen,  what  is  the  atomic  weight  of  iron  ? 


CHAPTER   XL 

EXPERIMENT  122. — Platinum. — Dip  some  asbestos  fiber 
into  a  solution  of  platinic  chloride,  PtCl4.  Use  the  solu- 
tion sparingly,  and  return  any  unused  portion  to  the 
bottle.  It  is  expensive.  Heat  the  asbestos  to  redness  for 
a  few  seconds,  holding  it  in  forceps.  The  result  of  this 
operation  is  to  coat  the  asbestos  with  finely  divided  plat- 


CARBON  131 

inum.    What  gas  must  have  escaped?    What  is  the  effect 
of  heat  upon  compounds  of  platinum? 

Let  the  "  platinized  asbestos,"  as  it  is  called,  cool,  hold 
it  in  forceps,  and  allow  hydrogen  from  a  jet  connected 
with  a  Kipp  generator  to  stream  out  against  it.  The 
generator  must  be  free  from  air,  or  an  explosion  will 
result.  If  there  is  any  doubt  about  this,  let  hydrogen 
escape  from  it  at  the  rate  of  a  bubble  a  second,  dipping 
the  exit-tube  in  water  for  ten  minutes  before  the  experi- 
ment is  tried.  Describe  and  explain  the  result. 


CHAPTER   XLI 

CARBON 

EXPERIMENT  123. — Allo tropic  forms  of  carbon. — Ex- 
amine graphite,  charcoal,  anthracite  and  bituminous  coal. 
Describe  them.  Describe  the  diamond.  Make  some  lamp- 
black by  holding  a  piece  of  chalk  in  a  luminous  gas  flame 
and  examine  and  describe  it. 

Color  some  water  faintly  with  litmus.  The  liquid 
must  be  clear  and  free  from  any  solid  matter.  Filter  it, 
if  necessary,  and  add  to  it  3-5  grams  of  animal  charcoal 
(bone-black).  Warm  the  liquid  gently,  shake  for  a  time, 
and  filter.  If  the  liquid  is  not  colorless,  filter  it  again. 
Repeat,  using  water  colored  with  a  little  ink  (one  drop) 
or  indigo,  instead  of  the  litmus.  This  remarkable  prop- 
erty of  decolorizing  liquids  is  also  possessed  by  wood  char- 
coal, but  to  a  far  less  degree. 

EXPERIMENT    124. — Illuminating    gas. — Fill    a    hard 

glass  test-tube  one-quarter  with  bituminous  coal,  clamp  ii 

horizontally  near  the  cork,  and  apply  a  heat,  at  first  gentle 

and  finally  intense.     During  the  heating  disconnect  sev- 

31 


132  ELEMENTARY  CHEMISTRY 

eral  times  and  test  the  gas  in  the  hard  glass  tube  for 
ammonia  with  red  litmus  paper.  Notice  the  production 
of  tar  during  the  heating.  Collect  the  gas  given  off  over 
water  in  wide  test-tubes.  Examine  it,  especially  with  re- 
gard to  combustibility.  Examine  the  residue  left  in  the 
tube,  which  is  coke. 


CHAPTER   XLII 

CARBON  DIOXIDE  AND    CARBON  MONOXIDE 

You  have  already  made  experiments  which  prove  that 
carbon  dioxide  is  produced  by  the  burning  of  charcoal  in 
the  air  or  in  oxygen,  and  that  it  is  contained  in  the  atmos- 
phere and  in  the  gases  from  the  lungs.  Bead  again  your 
notes  of  these  experiments,  and  if  you  have  forgotten 
any  essential  point,  repeat  them. 

EXPERIMENT  125. — Production  of  carbon  dioxide  by 
combustion. — Fill  a  deflagrating  spoon  with  powdered 
charcoal,  heat  it  to  redness,  and  let  it  burn  in  a  covered 
bottle.  Test  the  gas  with  lime-water. 

Hold  a  dry  clean  bottle  over  a  small  gas  flame  for  a 
few  seconds  and  apply  the  lime-water  test.  Repeat  with 
a  candle  flame.  With  burning  wood  (match-stick).  With 
the  flame  of  kerosene.  Of  alcohol.  (The  last  two  flames 
can  be  conveniently  obtained  by  dipping  a  bunch  of  as- 
bestos into  the  corresponding  liquid,  placing  it  on  an  iron 
plate  and  setting  fire  to  it. 

Draw  conclusions  regarding  the  existence  of  carbon 
in  combustibles  and  the  products  of  their  combustion. 
State  the  evidence. 

EXPERIMENT  126. — Production  of  carbon  dioxide  by 
the  action  of  acids  on  carbonates. — Place  a  little  sodium 


CARBON  DIOXIDE  AND  CARBON  MONOXIDE      133 

carbonate  in  each  of  three  small  beakers,  and  set  each  in 
a  larger  beaker  or  in  a  bottle  or  cylinder.  Add  to  the 
first  hydrochloric  acid,  to  the  second  nitric,  and  to  the 
third  sulphuric.  Cover  the  larger  vessels  with  paper. 
When  the  action  is  over,  remove  the  small  beakers  with 
forceps,  add  lime-water  to  each  of  the  larger  vessels,  and 
shake.  What  does  this  experiment  prove?  Write  the 
three  equations. 

EXPERIMENT  127. — Preparation  of  carbon  dioxide. — 
Use  the  apparatus  employed  in  making  hydrogen.  Place 
the  gas-bottle  almost  horizontal  and  slide  enough  broken 
marble  into  it  to  fill  it  when  upright  to  the  depth  of  1 
cm.  If  the  lumps  are  too  large  they  can  be  broken  with 
a  hammer  on  an  anvil — not  in  a  mortar.  Add  50  c.c.  of 
water,  and  then  hydrochloric  acid,  slowly  through  the 
funnel-tube  until  a  brisk  evolution  of  gas  is  obtained. 
Collect  it  by  downward  displacement  in  dry  bottles. 

Investigate  the  physical  properties  of  the  gas.  Has  it 
any  odor  or  taste?  If  in  doubt  about  the  last  point,  let 
the  gas  from  the  generator  bubble  through  a  little  water 
in  a  test-tube  and  taste  fhe  liquid.  Test  its  solubility  in 
water  in  the  usual  way. 

Lower  a  lighted  candle  into  a  bottle  of  the  gas.  In 
order  to  illustrate  its  high  density,  place  a  lighted  candle 
on  your  desk  and  pour  carbon  dioxide  over  it,  just  as 
you  would  pour  water.  Balance  a  large  beaker  on  the 
platform  scales  and  pour  a  bottle  of  the  gas  into  it. 

A  few  substances,  which  produce  very  high  tempera- 
tures by  their  combustion,  will  burn  in  the  gas.  Try  a 
piece  of  magnesium  ribbon,  held  in  forceps  and  started 
burning  in  the  air.  What  is  the  black  substance  which 
is  obtained  along  with  the  magnesium  oxide? 

Pass  carbon  dioxide  through  some  lime-water  in  a  test- 
tube  for  some  time.  Notice  that  the  calcium  carbonate 
which  it  at  first  precipitated  finally  redissolves.  Calcium 


134  ELEMENTARY  CHEMISTRY 

carbonate  is  soluble  in  water  containing  carbon  dioxide, 
and  this  solution  is  present  in  many  "hard  waters." 

Boil  the  liquid,  and  show  that  when  the  carbon  di- 
oxide is  expelled  the  precipitate  is  again  obtained. 

EXPERIMENT  128. — Action  of  glowing  charcoal  on  car- 
bon dioxide. — Pass  carbon  dioxide  from  a  Kipp  generator 
over  a  column  of  glowing  charcoal  about  20  cm.  long  in 
a  combustion-tube.  The  tube  is  heated  slowly  and  care- 
fully by  a  wing-top  burner.  Eed  rubber  stoppers  are  best 
for  it,  but  ordinary  corks  will  answer  with  care.  From 
this  tube  the  gas  passes  through  a  solution  of  potassium 
hydroxide  in  a  flask  with  a  doubly  perforated  cork.  It 
enters  through  a  tube  which  goes  almost  to  the  bottom 
of  the  flask,  and  leaves  through  a  tube  which  ends  just 
inside  the  cork  and  does  not  dip  into  the  liquid.  This 
tube  conveys  the  gas  to  a  pan,  where  it  is  collected  over 
water  in  wide  test-tubes.  The  current  of  gas  must  be  very 
slow. 

Study  the  properties  of  the  carbon  monoxide,  CO.  Es- 
pecially notice  its  combustion.  It  is  highly  poisonous, 
and  must  not  be  inhaled.  Apply  the  lime-water  test  to 
the  product  of  its  combustion. 

EXPERIMENT  129. — Preparation  of  carbon  monoxide 
by  heating  oxalic  acid  with  sulphuric  acid. — Place  8-10 
grams  of  oxalic  acid  and  50-70  grams  of  strong  sulphuric 
acid  in  a  small  flask.  Fit  the  flask  with  a  doubly  per- 
forated cork,  one  hole  of  which  carries  a  safety-funnel. 
A  little  mercury  must  be  placed  in  the  bend  of  the  safety- 
funnel.  Apply  a  very  gentle  heat.  The  gas  is  led  through 
a  wash-bottle  containing  potassium  hydroxide  and  col- 
lected over  water.  The  first  two  bottles  of  gas  contain 
air,  and  should  be  allowed  to  escape  under  the  hood. 
Investigate  the  properties  of  the  gas,  study  its  combustion 
by  burning  a  bottle  full  of  it,  and  test  the  product  of  the 
burning  by  lime-water. 


CARBON  DIOXIDE  AND  CARBON  MONOXIDE      135 

Care  must  be  taken  not  to  inhale  the  gas  produced  in 
this  experiment.  In  heating  the  flask,  use  a  small  flamef 
which  should  be  removed'  or  turned  down  if  the  contents 
begin  to  froth  up.  Keep  the  hand  away  from  under  the 
flask  in  heating,,  for  it  may  break,  and  hot  sulphuric  acid 
attacks  the  flesh  energetically. 

QUESTIONS 

1.  Why  is  carbon  monoxide  poisonous  t 

2.  Why  is  it  no  longer  poisonous  when  the  air  pressure  is 
greatly  increased  ? 

3.  What  volume  of  carbon  dioxide  will  be  produced  by  the 
burning  of  a  liter  of  carbon  monoxide  ?    Why  ?    What  volume  of 
oxygen  is  needed  ? 

PKOBLEMS 

147.  The  great  German  chemical  works,  the  Badische  Anilin- 
und  Soda  Fabrik,  burns  190,000  tons  of  coal  a  year.     If  the  coal 
contains  70  per  cent  of  carbon,  and  if  there  are  310  working  days 
in  the  year,  what  weight  of  COa  escapes  daily  from  the  chimneys 
of  the  establishment  ? 

148.  What  volume  of  carbon  dioxide  is  formed  by  the  burn- 
ing of  30  liters  of  carbon  monoxide,  and  what  volume  of  oxygen  is 
required  ?     Solve  by  inspection. 

149.  What  gas  gives  rise  to  the  blue  flame  often  seen  playing 
over  the  surface  of  a  coal  fire  ?    How  much  coal  containing  90  per 
cent  of  carbon  would  be  needed  to  make  5,000  liters  of  this  gas  at 
15°  and  750  mm.  ? 

150.  What  volume  of  gas  measured  under  standard  conditions 
is  produced  when  20  grams  of  pure  dry  oxalic  acid  are  heated  with 
strong  sulphuric  acid  ?    What  is  the  composition  of  the  gas  ? 

151.  What  volume  of  carbon  dioxide  must  be  passed  over 
glowing  charcoal  to  form  42  grams  of  carbon  monoxide  ? 

152.  What  is  the  volume  (a)  of  50  grams  of  carbon  monoxide  ? 
(5)  Of  50  grams  of  carbon  dioxide  ? 

153.  What  volume  of  carbon  dioxide  would  be  produced  by 
burning  a  diamond  weighing  3  grams  in  oxygen  ? 

154.  How  much  carbon  is  there  in  (a)  2.8  liters  of  carbon 
dioxide  ?     (&)  2.8  liters  of  carbon  monoxide  ? 


136  ELEMENTARY   CHEMISTRY 

155.  What  volume  of  carbon  dioxide  at  standard  conditions 
is  produced  by  dissolving  25  grams  of  marble  in  hydrochloric  acid  ? 

156.  What  volume  of  carbon  dioxide  at  12°  and  750  mm.  is 
produced  by  dissolving  20  grams  of  marble  in  hydrochloric  acid  ? 


CHAPTER   XLIII 

SOME   CARBON  COMPOUNDS 

EXPERIMENT  130. — Methane. — Heat  some  sodium  ace- 
tate in  an  iron  dish  until  the  water  of  crystallization  has 
escaped  and  the  melted  dry  salt  is  left.  Mix  3  grams  of 
this  in  a  mortar  with  3  grams  of  lime  and  3  grams  of 
sodium  hydroxide.  Grind  thoroughly  and  transfer  the 
mixture  to  a  hard  glass  test-tube.  Clamp  the  tube  hori- 
zontal and  fit  it  with  a  cork  bearing  a  delivery  tube. 
Apply  a  gentle  heat,  gradually  increased  to  redness.  Col- 
lect the  gas  in  wide  test-tubes.  Cover  the  hard  glass 
test-tube  with  soot,  and  disconnect  before  letting  it  cool. 

Eecord  the  physical  properties  of  the  gas.  Burn  some, 
notice  the  character  of  the  combustion,  and  apply  the 
lime-water  test. 

EXPERIMENT  131.— Acetylene.— Fill  a  small  test-tube 
one-third  with  water  and  throw  a  piece  of  calcium  car- 
bide into  it.  Has  the  gas  any  color  or  odor?  Light  it. 
Does  its  combustion  resemble  that  of  methane?  Eecord 
the  physical  properties  of  calcium  carbide. 

EXPERIMENT  132. — Various  carbon  compounds. — 
Ether,  wood  alcohol,  and  carbon  disulpliide  are  highly 
inflammable,  and  must  not  be  used  when  a  flame  is  any- 
where in  the  neighborhood.  The  bottles  containing  them 
must  be  kept  lightly  corked. 

Examine  the  following  important  carbon  compounds 
and  record  their  properties.  Study  also  their  solubility 


SOME  CARBON  COMPOUNDS  137 

in  water.  Alcohol,  ether,  chloroform,  wood  alcohol 
(methyl  alcohol),  and  carbon  disulphide.  Briefly  state 
in  your  notes  the  source  and  uses  of  each. 

EXPERIMENT  133.  —  Fermentation.  —  Dissolve  100 
grams  of  grape-sugar  (glucose)  in  a  liter  of  water.  Mix 
half  of  a  compressed  yeast  cake  to  a  thin  paste  with  water 
and  add  it  to  the  liquid.  Place  the  mixture  in  a  flask 
closed  by  a  singly  perforated  cork  bearing  a  tube  which 
leads  to  a  small  flask  containing  50  c.c.  of  lime-water. 
The  tube  should  dip  into  the  lime-water.  This  flask  is 
closed  by  a  doubly  perforated  cork,  and  a  tube  from  the 
second  hole  is  connected  with  a  U-tube  containing  pieces 
of  sodium  hydroxide.  This  is  to  prevent  the  C02  of  the 
air  from  entering  the  flask  containing  the  lime-water. 

Let  the  apparatus  stand  several  days,  if  possible. 
What  change  occurs  in  the  lime-water?  What  gas  must 
have  been  given  off?  What  is  the  other  product? 

To  obtain  it,  distill  the  liquid  in  the  large  flask  in  the  appara- 
tus employed  for  the  distillation  of  water  (Fig.  5).  Continue 
distilling  until  T\  of  the  liquid  has  passed  over,  and  discard  the 
rest.  If  there  is  too  much  liquid  to  be  distilled  in  one  operation, 
divide  it  into  several  portions,  and  distill  ^  of  each.  The  flask 
should  not  be  more  than  £  filled.  Unite  these  distillates  in  a 
small  flask,  and  distill  y1^  of  the  liquid,  using  a  dry  test-tube  as  a 
receiver.  Has  the  liquid  which  distills  the  odor  of  alcohol  ?  Is  it 
combustible  ?  If  not,  place  some  in  a  dish,  heat  it,  and  hold  a 
lighted  match  near  the  surface  of  the  liquid.1 

EXPERIMENT  134. — Aldehyde. — Make  a  strong  solu- 
tion of  potassium  dichromate,  place  it  in  a  test-tube,  and 
add  1  c.c.  of  strong  sulphuric  acid  and  1  c.c.  of  alcohol. 
Heat  the  liquid.  The  odor  produced  is  that  of  aldehyde. 
Notice  the  change  from  the  red  color  of  potassium  di- 

1  This  distillation  consumes  much  time  and  is  not  always  success- 
ful. It  is  sufficient  to  notice  the  production  of  C0a  and  the  general 
nature  of  the  process. 


138  ELEMENTARY  CHEMISTRY 

chromate  to  the  green  color  of  chromium  sulphate, 
Cr2(S04)3. 

EXPERIMENT  135. — Acetic  acid. — Examine  acetic  acid 
and  record  its  properties.  Dilute  some  with  100  times  its 
volume  of  water,  and  test  the  liquid  with  litmus  paper 
(red  and  blue).  Taste  it.  Mix  5  c.c.  of  strong  acetic 
acid  with  10  c.c.  of  alcohol.  If  possible,  preserve  some  of 
the  liquid  in  a  corked  test-tube,  and  notice  the  gradual 
development  of  the  fragrant  odor  of  acetic  ether.  This 
is  an  interesting  example  of  slow  chemical  change.  Sul- 
phuric acid  catalytically  accelerates  the  process.  Add  to 
some  of  the  fresh  mixture  in  a  test-tube  2  c.c.  of  strong 
sulphuric  acid  and  heat  gently.  The  odor  appears  at  once. 
This  is  a  test  for  acetic  acid  or  for  alcohol. 

Examine  vinegar,  taste  it,  and  try  its  behavior  with 
red  and  blue  litmus  paper. 

Examine  sodium  acetate,  and  notice  its  similarity  in 
general  character  to  ordinary  sodium  salts,  like  the  nitrate 
and  the  sulphate. 

QUESTIONS 

1.  What  is  fire-damp,  and  how  is  it  formed  ? 

2.  What  is  the  valence  of  carbon  in  CH4,  C2Hfl,  CsH8,  CHC18, 
CjH.Cl  ? 

3.  What  is  a  substitution  product  f    Give  examples.     Why  are 
substitution  products  so  numerous  ? 

4.  Write  the  formulas  of  the  first  nine  of  the  series  of  hydro- 
carbons which  begins  with  methane.     How  do  the  physical  proper- 
ties of  the  substances  change  as  we  advance  in  the  series  ? 

5.  Why  is  it  that  a  mixture  of  methane  or  any  other  combusti- 
ble gas  with  air  will  explode,  although  the  gas  alone  will  burn 
quietly  ? 

6.  What  is   fermentation  ?      State   briefly   two   methods  by 
which  it  has  been  proved  that  fermentation  is  not  caused  by  any 
vital  activity  of  yeast. 


ADDITIONAL  CARBON  COMPOUNDS  139 

PROBLEMS 

157.  What  is  the  weight  of  38  liters  of  methane  at  81*  ? 

158.  What  is  the  weight  of  10  liters  of  acetylene  ? 

159.  How  much  calcium  carbide  is  needed  to  produce  5.6 
liters  of  acetylene  ? 

160.  A  town  is  to  be  lighted  with  acetylene.     It  is  calculated 
that  the  consumption  of  the  gas  will  be  70,000  liters  per  day. 
How  much  calcium  carbide  will  be  required  per  month  of  thirty 
days? 

161.  How  much  carbon  is  there  in  32  liters  of  acetylene  ? 

162.  How  much  sodium  acetate  is  needed  to  make  8  liters  of 
methane  ? 

NaC,H,Oa  +  NaOH  =  CH4  +  NaaCO,. 

163.  How  many  liters  of  oxygen  at  10°  and  780  mm.  are 
needed  to  burn  completely  the  methane  obtained  when  41  grams 
of  sodium  acetate  are  heated  with  sodium  hydroxide  ? 

NaCaH3Oa  +  NaOH  =  CH4  +  NaaCO,. 
CH4  4-  203  =  CO,  +  2HaO. 

164.  Calcium  carbide  is  treated  with  water  and  the  acetylene 
burned.     28  liters  of  carbon  dioxide  at  15°  and  740  mm.  resulted 
from  the  combustion.     How  much  calcium  carbide  was  taken,  and 
what  volume  of  oxygen  at  15°  and  740  mm.  was  needed  to  burn 
the  acetylene  ? 

CaC2  +  2HaO  =  Ca(OH)a  +  C2H,. 
C9H2  +  5O  =  2C02  +  HaO. 


CHAPTEE   XLIV 

ADDITIONAL    CARBON  COMPOUNDS 

EXPERIMENT  136. — Soap. — Dissolve  25  grams  of  so- 
dium hydroxide  in  150  c.c.  of  water  in  an  iron  dish.  Add 
75  grams  of  lard,  and  boil  gently  for  half  an  hour  or 
more.  Slowly  add  50  grams  of  salt,  stirring  constantly. 
Be  careful  during  the  boiling  not  to  allow  any  of  the 


140  ELEMENTARY  CHEMISTRY 

liquid  to  be  spattered  into  the  eyes.  It  is  well  to  keep 
a  glass  plate  between  the  dish  and  the  face. 

The  solid  which  separates  at  the  top  is  soap.  What 
else  has  been  produced,  and  where  is  it?  Kemove  the 
soap  and  examine  it.  It  is  a  mixture  of  the  sodium  salts 
of  palmitic,  stearic,  and  oleic  acids.  Dissolve  about  3 
grams  of  shavings  of  it  in  100  c.c.  of  warm  water  in  a 
dish.  Add  dilute  sulphuric  acid  to  a  portion  of  it.  The 
mass  which  separates  consists  of  a  mixture  of  the  acids 
mentioned  above,  which  are  liberated  from  their  sodium 
salts  by  the  sulphuric  acid,  sodium  sulphate  being  formed. 
To  another  small  portion  of  the  soap  solution  add  a  dilute 
solution  of  calcium  chloride.  Water  containing  calcium 
compounds  is  said  to  be  hard.  Why  can  not  hard  water 
be  used  satisfactorily  in  washing  with  soap? 

EXPERIMENT  137. — Albumin. — Break  an  egg  and  sepa- 
rate some  of  the  white  from  the  yolk.  Dilute  about  5  c.c. 
of  the  white  of  egg  with  100  c.c.  of  water.  Boil  a  portion 
of  this  liquid  in  a  test-tube.  (?)  To  another  portion  in 
a  test-tube  add  a  little  nitric  acid.  The  albumin  is  pre- 
cipitated in  both  experiments. 

Heat  a  bit  of  meat,  a  feather,  a  few  clippings  of  horn, 
or  a  cochineal  insect  in  a  dry  test-tube.  Almost  any  form 
of  animal  matter  will  answer.  Describe  the  nature  of  the 
change.  Notice  the  odor.  Show  by  red  litmus  paper  that 
ammonia  is  given  off.  This  proves  that  nitrogen  and  hy- 
drogen are  two  of  the  constituents  of  albumin.  Look  for 
evidence  of  the  presence  of  carbon  in  the  residue  in  the 
tube. 


CHAPTEE   XLV 

No  experiments. 


ADDITIONAL  CARBON  COMPOUNDS  141 

CHAPTER   XLVI 

No  experiments. 

Additional  exercise.  The  fundamental  principle!  of 
qualitative  analysis.1 — A.  Prepare  dilute  solutions  of  the 
nitrates  of  silver,  copper,  iron,  barium,  and  sodium. 
About  100  c.c.  of  each  will  be  needed.  Copper  nitrate 
can  be  made  by  heating  a  little  copper  in  a  test-tube  with 
dilute  nitric  acid  until  the  action  ceases.  The  liquid 
should  still  contain  undissolved  copper.  It  is  diluted  and 
filtered  if  necessary,  or  simply  poured  off  from  the  cop- 
per. Iron  nitrate  (ferric)  can  be  made  in  a  similar  way, 
but  the  iron  must  be  well  covered  with  water  and  the 
nitric  acid  added  drop  by  drop,  for  the  action  is  violent. 
The  other  solutions  can  be  made  by  dissolving  1  gram  of 
the  corresponding  nitrate  in  100  c.c.  of  water. 

Study  the  behavior  of  small  portions  of  each  of  these 
liquids  with  the  following  substances  separately: 

B.  Hydrochloric  acid,  sulphuric  acid,  hydrogen  sul- 
phide, ammonium  hydroxide. 

The  experiments  should  be  made  in  test-tubes.  The 
liquids  in  B  should  be  added  one  drop  at  a  time,  stirring 
constantly.  Note  especially  whether  a  precipitate  or  a 
color  is  produced,  and  if  the  former,  whether  it  is  soluble 
in  excess.  The  hydrogen  sulphide  can  be  taken  from  a 
Kipp  generator.  Use  a  gentle  current  of  the  gas,  not 
more  than  a  bubble  every  two  seconds.  The  tube  through 
which  the  H28  passes  into  the  liquid  must  be  cleaned 

1  The  object  of  this  exercise  is  not  to  teach  qualitative  analysis, 
but  to  give  the  student  a  grasp  of  the  principles  upon  which  the 
separation  of  one  metal  from  another  depends,  so  that  he  will  be 
able  afterward  to  attack  the  subject  understandingly.  The  scheme 
is  that  of  Professor  Richards  (Harvard  Requirements,  p.  20). 


142  ELEMENTARY  CHEMISTRY 

both  from  liquid  and  solid  before  using  it  again  for  a 
fresh  test,  and  every  test-tube  and  beaker  employed  must 
be  absolutely  clean.  If  no  Kipp  generator  is  available, 
make  the  hydrogen  sulphide  yourself  by  the  method  al- 
ready studied.  Be  careful  not  to  inhale  it  unnecessarily. 
NOTE. — The  precipitate  formed  when  the  solution  of 
ferric  nitrate  is  treated  with  H2S  contains  no  iron.  It  is 
sulphur.  The  change  in  the  color  of  the  liquid  will  show 
you  that  the  iron  is  now  in  the  ferrous  condition. 

QUESTIONS 

1.  Devise  a  method  of  separating  silver  from  any  of  the  metals 
in  A  based  on  the  use  of  hydrochloric  acid.     Try  it,  mixing  a  little 
of  the  silver-nitrate  and  copper- nitrate  solutions  for  the  purpose. 
Why  is  it  necessary  to  wash  the  precipitate  by  filling  the  filter  with 
water  and  letting  it  drain  off  half  a  dozen  times  after  filtering  ? 
This  must  always  be  done  in  work  of  this  kind.     In  the  present 
exercise  the  wash  liquid  can  be  thrown  away.     Of  course,  if  the 
work  is  quantitative  it  must  be  retained.     Why  is  it  necessary  to 
add  the  hydrochloric  acid  until  it  produces  no  further  precipitate  ? 
This  is  always  necessary  in  similar  cases.     How  can  you  prove  that 
the  precipitate  is  really  silver  chloride  ?     How  could  you  tell  after 
filtering  whether  the  silver  had  all  been  precipitated  or  not  ? 

2.  Devise  a  method  of  separating  barium  from  any  of  the 
others.     Try  it  with  barium  and  copper  nitrates. 

3.  Devise  and  carry  out  a  method  of  separating  silver,  copper, 
and  barium,  when  all  three  are  present  together. 

4.  Devise  and  carry  out  a  method  of  separating  copper  from 
sodium  by  hydrogen    sulphide.      Evaporate  the   liquid,    which 
should  contain  the  sodium,  to  dryness,  and  apply  the  flame  test. 
Identify  the  copper  as  in  5  &  (next  paragraph). 

5.  Separate  copper  from  iron  by  hydrogen  sulphide,     (a)  Heat 
the  filtrate  containing  the  ferrous  ions  to  gentle  boiling  for  twenty 
minutes  to  drive  off  the  H2S.     Add  a  few  drops  of  nitric  acid,  and 
boil.     What  change  occurs  in  the  condition  of  the  iron  ?     Show 
that  the  liquid  contains  iron,  using  ammonium  hydroxide.     (6) 
Dissolve  the  copper  sulphide  by  repeatedly  pouring  the  same  por- 
tion of  hot  dilute  nitric  acid  through  the  filter.     The  blue  liquid 


ADDITIONAL  CARBON  COMPOUNDS  143 

contains  copper  nitrate.     Refer  to  your  previous  notes  for  several 
methods  of  identifying  copper  in  it,  and  apply  them. 

6.  Separate  iron  from  sodium,  using  ammonium  hydroxide. 

7.  Devise  and  carry  out  a  method  of  separating  copper,  iron, 
and  sodium  when  all  three  are  present  together. 

8.  Devise  and  carry  out  a  method  for  the  separation  of  all  five 
of  the  metals  in  A,  using  only  the  substances  in  B  for  the  separa- 
tion.    Any  method  you  are  familiar  with  can  be  used  to  identify 
the  metals  when  separated, 


APPENDIX 


CALCULATION  OF  THE  EFFECT  OF  TEMPERATURE, 
PRESSURE,  AND  WATER-VAPOR  ON  THE  VOLUMES 
OF  GASES 

1.  Temperature. — The  absolute  temperature  is  the  tem- 
perature measured  from  273°  below  0°  C.  Thus,  the 
absolute  temperature  of  10°  is  273  +  10  =  283°. 

The  absolute  temperature  of  —10°  is  273-10  =  263°. 

PROBLEM 

165.  Calculate  the  absolute  temperature  corresponding  to  the 
following  centigrade  temperatures  : 

a.  13°.  b.  274°.  c.  -  50°.  d.  -  273°. 

The  volume  of  a  mass  of  gas  is  directly  proportional 
to  its  absolute  temperature.  Let  T  and  t  =  two  tempera- 
tures— both  absolute.  Let  VT  be  the  volume  at  T°  and  V 
be  the  volume  at  t°.  Then — 

T 
FT  =  F,xf 

V 

In  words,  this  means  if  you  know  the  volume  of  a 
mass  of  gas  at  some  known  temperature,  you  can  calculate 
its  volume  at  some  other  temperature  by  multiplying  the 
old  volume  by  the  new  temperature,  and  dividing  by  the 
old  temperature,  both  temperatures  being  absolute. 

Never  make  the  error  of  using  ordinary  centigrade 
temperatures  instead  of  absolute  temperatures.  It  is  easy 
to  see  that  this  leads  directly  to  absurd  results.  For  in- 

145 


146  ELEMENTARY  CHEMISTRY 

stance:  I  have  a  liter  of  gas  at  0°  C.    What  will  its  volume 
become  at  273°  C.  ? 

Here,  if  we  use  ordinary  centigrade  degrees,  the  vol- 
ume becomes 

273 

ix-5-; 

that  is,  the  volume  is  infinite — which  is  absurd.     But  if 
we  employ  absolute  degrees,  the  volume  is — 

IX  fff  =  8  liters; 
which  is  the  correct  result. 

PROBLEMS 

166.  What  volume  will  a  liter  of  air  at  0°  C.  occupy  at  100°  C.  ? 

167.  5  liters  of  oxygen  at  0°  C.  occupy  what  volume  (a)  at 
10°  C.  ?     (&)  at  —10°  C.  ? 

168.  25  c.c.  of  nitrogen  at  15°  C.  will  measure  what  at  the 
standard  temperature  0°  C.  ? 

169.  I  have  500  c.c.  of  hydrogen  at  13°  C.     What  will  the 
volume  become  at  65°  C.  ? 

170.  600   c.c.    of  oxygen   at   28°   C.   will  measure   what  at 
-14°  C.  ? 

171.  500  liters  of  air  at  20°  C.  will  occupy  what  volume  at 
80°  C.? 

172.  A  liter  of  steam  at  100°  C.  will  occupy  what  volume  at 
120°  C.? 

173.  67  liters  of  air  are  heated  from  —30°  C.  to  60°  C.  ?    What 
is  the  new  volume  ? 

Since  the  volume  of  a  mass  of  gas  varies  with  the 
temperature,  it  is  always  necessary,  in  measuring  gases, 
to  know  the  temperature  of  the  gas  measured.  And  it 
is  clear  that  the  expression  "  1  liter  of  oxygen  "  has  no 
meaning  unless  some  particular  temperature  is  either 
stated  or  understood.  Now,  in  order  to  avoid  the  neces- 
sity of  continually  stating  the  temperature,  it  is  extremely 
convenient  to  assume  some  temperature  as  a  standard 


APPENDIX  147 

point  which  is  to  be  understood  unless  some  other  tem- 
perature is  stated.  The  standard  temperature  universally 
agreed  upon  is  0°  C.  —  the  melting-point  of  ice.  Thus, 
when  a  writer  speaks  of  "  1  liter  of  oxygen  "  without 
stating  the  temperature  under  which  the  gas  was  meas- 
ured, we  know  that  0°  C.  is  meant. 

The  student  should  grasp  the  fact  that  every  problem 
like  those  just  solved  is  supposed  to  deal  with  a  certain 
fixed  weight  of  gas  which  is  not  added  to  or  subtracted 
from  during  the  process  of  heating  or  cooling.  Clearly, 
if-  temperature  and  pressure  remain  the  same,  the  volume 
must  be  directly  proportional  to  the  weight  of  the  gas. 
Thus,  1  gram  of  hydrogen  at  standard  temperature  and 
pressure  occupies  a  volume  of  11.2  liters.  Evidently,  2 
grams  of  hydrogen  must  measure  22.4  liters  under  the 
same  conditions,  and  so  on.  But,  in  all  problems  of  this 
sort,  the  quantity  of  gas  is  supposed  to  remain  the  same. 

2.  Pressure.  —  The  volume  of  a  mass  of  gas  is  inversely 
proportional  to  the  pressure  upon  it.  If  p  and  P  are 
two  pressure,  both  stated  in  millimeters  of  mercury,  and 
if  Vp  and  FP  are  the  volumes  which  the  same  quantity 
of  gas  will  occupy  at  those  pressures,  then 

FP  :  Vp  ::  p  :  P.    Hence 


Therefore,  if  the  volume  of  a  mass  of  gas  is  given  at 
some  pressure  and  it  is  required  to  calculate  its  volume 
at  some  other  pressure,  we  must  multiply  the  old  volume 
by  the  old  pressure  and  divide  by  the  new  pressure. 

PROBLEMS 

174.  10  liters  of  gas  at  a  pressure  of  743  mm.  will  occupy 
what  volume  at  720  mm.  ? 

175.  18.5  c.c.  of  nitrogen  are  measured  under  a  pressure  of 
745  mm.     What  will  the  volume  be  at  760  mm.  ? 


148  ELEMENTARY  CHEMISTRY 

176.  A  liter  of  oxygen  is  measured  at  760  mm.     What  will  it 
measure  at  748  mm.  ? 

177.  100  c.c.  of  air  at  760  mm.  (1  atmosphere)  will  occupy 
what  volume  under  20  atmospheres  ? 

178.  What  pressure  is  required  to  compress  500  c.c.  of  carbon 
dioxide  at  728  mm.  to  a  volume  of  400  c.c.  ? 

179.  What  must  the  pressure  be  made  in  order  to  allow  the 
500  c.c.  of  gas  of  the  preceding  problem  to  expand  to  850  c.c.  ? 

In  order  not  to  be  compelled  to  state  continually  the 
pressure,  in  speaking  of  the  volumes  of  gases,  and  in  order 
to  be  able  to  compare  gas  volumes,  measured  at  different 
temperatures,  with  each  other,  760  mm.  of  mercury  is 
agreed  upon  as  the  standard  pressure,  which  is  under- 
stood when  no  pressure  is  stated.  This  pressure  is  called 
1  atmosphere,  because  the  pressure  of  the  air  does  not  vary 
widely  from  that  amount. 

Since,  as  we  have  seen,  0°  is  the  standard  temperature, 
the  expression  "  standard  conditions  "  means  0°  and  760 
mm.  Thus,  when  a  writer  speaks  of  1  liter  of  oxygen  (or 
of  any  volume  of  any  gas)  without  mentioning  either  tem- 
perature or  pressure,  we  understand  at  once  that  the  gas  is 
supposed  to  exist  at  0°,  and  under  a  pressure  of  760  mm. 

3.  When  temperature  and  pressure  both  vary,  we  have 
simply  to  correct  for  both  by  the  methods  already  studied. 
This  can  be  done  in  two  separate  calculations,  but  it  is 
easier  and  better  to  unite  both  corrections  in  one  opera- 
tion. The  volume  of  a  gas  is  directly  proportional  to  the 
absolute  temperature  and  inversely  proportional  to  the 
pressure.  Let 

Fpr  =  the  volume  at  the  absolute  temperature   T  and 

pressure  P. 
Vft  =  the  volume    at   the   absolute    temperature   t  and 

pressure  p. 

Then  v  T      p 

t       P 


APPENDIX  149 

In  words,  this  means  that  in  order  to  calculate  the 
new  volume  of  a  gas  at  some  new  temperature  and  pres- 
sure, we  must  multiply  the  old  volume  by  the  new  tem- 
perature and  the  old  pressure,  and  divide  it  by  the  old 
temperature  and  the  new  pressure.  Of  course,  both  tem- 
peratures must  be  absolute. 

Such  calculations  can  be  rapidly,  easily,  and  correctly 
made  by  the  use  of  logarithms,  and  this  is  true  of  chem- 
ical calculations  generally.  A  table  of  logarithms  is  given 
for  this  purpose,  and  its  use  will  save  about  half  the  time 
and  labor  of  chemical  calculation,  and  will  greatly  reduce 
the  number  of  errors  in  the  numerical  work. 

PKOBLEMS 

180.  100  c.c.  of  oxygen  at  15°  C.  and  740  mm.  will  occupy 
what  volume  at  standard  conditions  ? 


The  student  will  find  that  his  chief  difficulty  in  solv- 
ing problems  like  this  and  the  following  ones  is  in  de- 
termining which  temperature  and  pressure  to  put  in  the 
numerator  and  which  in  the  denominator.  It  will  pay  to 
make  it  a  rule  to  inspect  the  fractions  with  great  care 
before  working  out  the  calculation.  Errors  can  be  de- 
tected by  the  exercise  of  a  little  common  sense.  For  in- 
stance, in  the  preceding  problem  the  gas  is  to  be  cooled 
from  15°  C.  to  0°  C.  This  will  reduce  its  volume.  Hence, 


288 
the  temperature-fraction  must  be        ,  not        .    Also,  the 


pressure  is  to  be  raised  from  740  to  760,  and  this  also  will 
reduce  the  volume.     Hence,  the  pressure  fraction  must 

740       ,   760 
be760'not  740' 

181.  Supposing  the  initial  temperature  in  the  preceding  prob- 
lem to  be  —15°  C.  instead  of  15°  C.,  what  would  be  the  new  vol- 
ume ?  The  other  figures  remain  the  same. 


150  ELEMENTARY  CHEMISTRY 

182.  What  volume  will  48  c.c.  of  nitrogen  at  standard  condi- 
tions occupy  at  18°  C.  and  733  mm.  ? 

183.  25  liters  of  a  gas  at  standard  conditions  are  cooled  to  — 10° 
C.,  and  the  pressure  reduced  to  723  mm.    What  is  the  new  volume  ? 

184.  310  c.c.  of  hydrogen  at  10°  C.  and  530  mm.  will  occupy 
what  volume  at  18.7°  C.  and  590  mm.  ? 

185.  1,704  c.c.  of  nitrogen  at  11°  C.  and  760  mm.  are  brought 
to  a  temperature  of  27°  C.  and  a  pressure  of  900  mm.     What  is  the 
volume  ? 

186.  271  c.c.  of  hydrogen  at  269°  C.  and  900  mm.  are  cooled 
to  — 51°  C.,  and  the  pressure  decreased  to  666  mm.     Calculate  the 
final  volume. 

4.  The  effect  of  water-vapor  on  the  volume  of  a  mass 
of  gas. — Suppose  that  we  have  100  c.c.  of  dry  oxygen  con- 
fined over  mercury  in  a  graduated  tube.  Let  us  admit 
a  drop  of  water  and  allow  the  oxygen  to  saturate  itself 
with  moisture.  Clearly,  the  volume  of  gas  in  the  tube 
must  increase,  for  the  water-vapor  will  occupy  space.  The 
result  is  the  same  as  though  we  had  introduced  a  little 
nitrogen  or  some  other  gas  into  the  tube  and  allowed  it 
to  mix  with  the  oxygen. 

The  volume  can  be  kept  100  c.c.  by  increasing  the 
pressure  under  which  the  gas  is  measured.  But  if  this 
is  done,  the  total  pressure  can  not  be  considered  as  ex- 
erted upon  the  oxygen  in  the  tube,  for  the  water-vapor 
is  also  present.  Hence,  the  pressure  under  which  the  gas 
really  exists  and  is  measured  is  less  than  the  total  pressure. 
How  much  less? 

The  pressure  which  saturated  water-vapor  exerts  at 
various  temperatures  is  given  in  the  table.  When  a  gas 
is  measured  over  water,  or  when  it  is  measured  saturated 
with  water,  the  pressure  which  water-vapor  exerts  at  the 
temperature  of  measurement  must  be  ascertained  from  the 
table  and  deducted  from  the  total  pressure.  The  remain- 
der will  be  the  pressure  under  which  the  gas  is  really 
measured. 


APPENDIX  151 

PKOBLEMS 

187.  A  mass  of  air  at  15.3°  C.  and  747.2  mm.,  measured  over 
water,  occupied  a  volume  of  82.4  c.c.  What  volume  would  it 
occupy  dry  and  at  standard  conditions  ? 

SOLUTION 

From  the  table  we  observe  that  water-vapor  at  15°  C.  exerts 
a  pressure  of  12.7  mm.  and  at  16°  C.  a  pressure  of  13.54  mm. 
Hence  its  pressure  at  15.3°  C.  must  =  12.9  mm. 

The  pressure  under  which  the  gas  is  really  measured  is 

747.2  —  12.9  =  734.3  mm. 

The  rest  of  the  calculation  is  the  same  as  in  the  preceding 
problems  : 


188.  11.41  c.c.   of  a  mixture  of  oxygen  and  hydrogen  are 
measured  over  water  at  14°  C.  and  743  mm.     Calculate  the  volume 
under  standard  conditions. 

189.  112.1  c.c.  of  nitrogen  saturated  with  water  at  16°  C.  and 
744  mm.  will  occupy  what  volume  dry  and  under  standard  condi- 
tions ? 

190.  The  gas-holder  of   a  gas-works  contains  4,500  cubic 
meters  of  illuminating  gas,  confined  over  water.     The  temperature 
is  9*  C.  and  the  pressure  776  mm.     How  many  cubic  meters  would 
the  gas  measure  under  standard  conditions  ? 

191.  100  c.c.  of  oxygen  are  confined  over  water  and  measured 
at  14°  C.  and  756  mm.     What  will  be  the  volume  when  the  gas  is 
dried  and  placed  under  standard  conditions  ? 

192.  A  gas-holder  contains  10  liters  of  air  confined  over  water 
at  20?  C.  and  756  mm.     What  will  the  gas  measure  when  dried, 
other  conditions  remaining  the  same  ? 


152 


ELEMENTARY  CHEMISTRY 


TABLE  OF  ATOMIC  WEIGHTS 


NAME. 

Symbola. 

Exact 

values. 

| 

|j 

S  ^ 

•< 

NAME. 

Symboli. 

Exact 

values. 

Approximate 
value..  1 

Aluminium  

A.1 

27.1 

27 

Neodymium  

Nd 

143.6 

Sb 

120  0 

120 

Neon 

NP 

19.94 

A 

3992 

Nickel  . 

Ni 

58.70 

587 

Arsenic  

As 
Rn 

75.0 
137  43 

75 
137 

••Nitrogen  
Osmium 

N 
O 

14.04 
190.8 

14 

Beryllium  

Be 

9.1 

Oxygen  

O 

16.000 

16 

Bismuth 

Bi 

208  0 

208 

Palladium 

Pd 

106.5 

P 

11  0 

H 

>  Phosphorus 

p 

31.0 

31 

Br 

79  9-"5 

80 

Platinum 

Pt 

195.2 

195 

Cadmium 

Od 

112  3 

112 

*^otassium     .  . 

K 

39.14 

39 

Caesium  

f!s 

1329 

Praseodymium  .  .  . 

Pr 

140.5 

v  Calcium. 

Pfl 

40  1 

40 

Radium 

Rfl 

225.? 

"  Carbon 

p 

12  001 

12 

Rhodium 

P,h 

103.0 

Cerium 

rp 

140  0 

Rubidium  

Pb 

85.44 

*  Chlorine.         

ci 

35  455 

35  5 

Ruthenium  . 

Pitl 

101  7 

Chromium  

Or 

52  14 

52 

Samarium 

Sfl 

1500 

"Cobalt  

Po 

59  00 

59 

Scandium 

So 

440 

Columbium  

rb 

94  0 

Se 

792 

•'  Copper 

Cu 

63  60 

63  5 

Si 

28  4 

28  5 

Erbium  

Fr 

166  0 

Silver 

\rr 

10?  93 

108 

Fluorine  

F 

19.05 

19 

vSodium  

Na 

23.05 

23 

Gadolinium  
Gallium  .         

Gd 
Ga 

156.  ? 
70  0 

Strontium  

Sr 

g 

87.68 
32065 

87.5 
32 

Germanium  

Ge 

72  5 

Tantalum 

TH 

183.0 

Gold 

\u 

•107  q 

1Q7 

TP 

127  5  * 

/Helium  

Ho 

3  96 

Tb 

160  0 

*  Hydrogen  

H 

1  0075 

1 

Tl 

204  15 

Indium 

In 

Th 

003  n 

/Iodine  ,  

T 

126  85 

127 

Tu 

171  0  ? 

Iridium  

Tr 

1930 

Tin 

8n 

1190 

119 

/  Iron  

Fe 

55  9 

56 

Ti 

48  17 

Krypton  

Kr 

81  7 

184  0 

Lanthanum  

T.fl 

138  5 

u 

238  5 

239 

^Lead  

Ph 

206  92 

207 

•y' 

fit  4 

r  Lithium  

Tii 

7.03 

x 

1280 

„  Magnesium..,  

Mf 

24  36 

24 

Ytterbium 

Yb 

173  0 

f  Manganese 

Mn 

55  02 

•f  Mercury  . 

He 

200  0 

cr  e 

Molybdenum  

Mo 

96.0 

Zirconium  

Zr 

90.6 

This  table  contains  the  values  of  the  atomic  weights 
according  to  the  calculations  of  Prof.  Richards.  The 
(ij'/iroximate  values  should  be  used  in  solving  all  problems. 


APPENDIX 


153 


VAPOR  PRESSURE  OF  WATER 


Tempera- 
ture, Centi- 
grade. 

Vapor 
pressure  in 
mm.  of 
mercury. 

Tempera- 
ture, Cen- 
tigrade. 

Vapor 
pressure  in 
mm.  of 
mercury. 

Tempera- 
ture, Cen 
tigrade. 

Vapor  pressure 
in  mm.  of 
mercury. 

-10 

2.09 

12 

10.46 

26 

24.99 

-  5 

3.11 

13 

11.16 

27 

26.51 

0 

4.60 

14 

11.91 

28 

28.10 

+   1 

4.94 

15 

12.70 

29 

29.78 

2 

5.30 

16 

13.54 

30 

31.55-k 

3 

5.69 

17 

14.42 

35 

41.  83  I 

4 

6.10 

18 

15.36 

40 

54.91 

5 

6.53 

19 

16.35 

50 

91.98 

6 

7.00 

20 

17.39 

60 

148.79 

7 

7.49 

21 

18.50 

70 

233.09 

8 

8.02 

22 

19.66 

80 

354.64 

9 

8.57 

23 

20.89 

90 

525.45 

10 

9.17 

24 

22.18 

100 

760.00 

11 

9.79 

25 

23.55 

The  vapor  pressure  of  water  for  a  temperature  not 
given  in  the ,  table  can  easily  be  found  by  calculation. 
Thus,  suppose  it  is  required  to  find  the  vapor  pressure  for 
the  temperature  of  32.5°.  The  increase  in  vapor  pressure 
from  30°  to  35°  is  41.83  —  31.55  =  10.28  mm.  Hence, 
the  increase  from  30°  to  32.5°  will  not  be  far  from 

10.28  X  — ,  or   5.14  mm.,  and  the   vapor   pressure   for 

32.5°  will  be  about  36.69  mm.  It  will  not  be  exactly 
36.63  mm.,  because,  in  the  calculation,  it  is  assumed  that 
the  vapor  pressure  increases  proportionally  with  the  tem- 
perature, which  is  not  the  case,  but  for  small  differences 
of  temperature  the  error  is  small. 

When  a  gas  is  measured  over  water,  or  moist,  the  vapor 
pressure  of  water  for  the  temperature  must  be  subtracted 
from  the  pressure  under  which  the  gas  is  measured. 


154  ELEMENTARY   CHEMISTRY 

FACTORS  FOB  CONVERTING  METRIC  INTO  ORDINARY 
UNITS 

1  inch  =  2.54  cm.  1  cm.  =  0.3937  in. 

For  practical  purposes  it  is  sufficient  to  remember  that 
about  2%  cm.  =  1  in. 

A  liter  is  the  volume  of  a  cube  whose  side  is  10  cm. 
Therefore, 

1  liter     =  1,000  cubic  centimeters ; 
1  pint     =  0.47318  liter; 
1  gallon  =  3.78543  liters; 
1  liter     =  0.26417  gallon. 

The  gram  is  the  weight  of  1  c.c.  of  pure  water  at  4°. 
1  liter  of  pure  water  at  4°  =  1  kilo  (1,000  grams). 
1  oz.  =  28.35  gms.  1  gram  =  15.432  grains. 

1  Ib.  =  453.6  gms.  1  kilo    =  2.2046  Ibs. 

FORMULA  FOR  CONVERTING  FAHRENHEIT  DEGREES 
INTO  CENTIGRADE,  AND  THE  REVERSE 

C.°  =  ^(F.°—  32). 
y 

F.°  =  |  C.°  +  32. 

MISCELLANEOUS  DATA 

Weight,  1  liter  pure  hydrogen  at  0°  and  760  mm.  = 
.0896  gram.  Volume  of  molecular  weight  in  grams  of 
any  gas  or  vapor,  22.4  liters. 

Melting-point  of  ice 0°  C, 

"  alcohol -130°  C. 

"  "   mercury — 39.4°  C. 

"  "  hydrogen,  about —260°  C. 

"  "  tin 233°  C. 

"          "  lead 334°  C. 

"  "   silver 954°  C. 

"  gold 1,064°  C. 

"  "  iron  (pure) 1,600°  C. 


APPENDIX  155 

Boiling-point  of  water 100°  C. 

"  "   alcohol 78°  C. 

"  "   ether 35°  C. 

"  "   chloroform 61.5°  C. 

"  "   carbon  disulphide. . .  47°  C. 

"  "   hydrogen —  252.8°  C. 

«          "  oxygen -183°  C. 

Air  contains  21  per  cent  (20.97  per  cent)  by  volume  of 
oxygen.  1  liter  of  dry  air  at  0°  and  760  mm.  weighs  1.293 
grams. 


Logarithms. 


Proportional  parts 


!Nat. 
Kumbtr 

0 

1 

2 

3 

4 

5 

6 

7 

8 

9 

1  2  3 

456 

789 

10 

0000 

0043 

0086 

0128 

0170 

0212 

0253 

0294 

0334 

0374 

4  8  12 

17  21  25 

29  33  37 

11 

0414 

0453 

0492 

0531 

0569 

0607 

0645 

0682 

0719 

0755 

4  8  11 

15  19  23 

26  30  34 

12 

0792 

0828 

0864 

0899 

0934 

0969 

1004 

1038 

1072 

1106 

3  7  10 

14  17  21 

24  28  31 

13 

1139 

1173 

1206 

1239 

1271 

1303 

1335 

1367 

1399 

1430 

3  6  10 

13  16  19 

23  26  29 

14 

1461 

1492 

1523 

1553 

1584 

1614 

1644 

1673 

1703 

3732 

3  6  9 

12  15  18 

21  24  27 

15 

1761 

1790 

1818 

1847 

1875 

1903 

1931 

1959 

1987 

2014 

368 

11  14  17 

20  22  25 

16 

2041 

2068 

2095 

2122 

2148 

2175 

2201 

2227 

2253 

2279 

358 

11  13  16 

18  21  24 

17 

2304 

2330 

2355 

2380 

2405 

2430 

2455 

2430 

2504 

2529 

257 

10  12  15 

17  20  22 

18 

2553 

2577 

2601 

2625 

2648 

2672 

2695 

2718 

2742 

2765 

257 

9  12  14 

16  19  21 

19 

2788 

2810 

2833 

2856 

2878 

2900 

2923 

2945 

2967 

2989 

247 

9  11  13 

16  18  20 

20 

3010 

3032 

3054 

3075 

3096 

3118 

3139 

3160 

3181 

3201 

246 

8  11  13 

15  17  19 

21 

3222 

3243 

3263 

3284 

3304 

3324 

3345 

3365 

3385 

3404 

246 

8  10  12 

14  16  18 

22 

3424 

3444 

3464 

3483 

3502 

3522 

3541 

3560 

3579 

3598 

246 

8  10  12 

14  15  17 

23 

3617 

3636 

3655 

3674 

3692 

3711 

3729 

3747 

3766 

3784 

246 

7  9  11 

13  15  17 

24 

3802 

3820 

3838 

3856 

3874 

3892 

3909 

3927 

3945 

3962 

245 

7  9  11 

12  14  16 

25 

3979 

3997 

4014 

4031 

4048 

4065 

4082 

4099 

4116 

4133 

235 

7  9  10 

12  14  15 

26 

4150 

4166 

4183 

4200 

4216 

4232 

4249 

4265 

4281 

4298 

235 

7  8  10 

11  13  15 

27 

4314 

4330 

4346 

4362 

4378 

4393 

4409 

4425 

4440 

4456 

235 

689 

11  13  14 

28 

4472 

4487 

4502 

4518 

4533 

4548 

4564 

4579 

4594 

4609 

235 

689 

11  12  14 

29 

4624 

4639 

4654 

4669 

4683 

4698 

4713 

4728 

4742 

4757 

134 

679 

10  12  13 

30 

4771 

4786 

4800 

4814 

4829 

4843 

4857 

4871 

4886 

4900 

1  3 

679 

10  11  13 

31 

4914 

4928 

4942 

4955 

4969 

4983 

4997 

5011 

5024 

5038 

3 

678 

10  11  12 

32 

5051 

5065 

5079 

5092 

5105 

5119 

5132 

5145 

5159 

5172 

3 

578 

9  11  12 

33 

5185 

5198 

5211 

5224 

5237 

5250 

5263 

5276 

5289 

5302 

3 

5  6  8 

9  10  12 

34 

5315 

5328 

5340 

5353 

5366 

5378 

5391 

5403 

5416 

5428 

3 

568 

9  10  11 

35 

5441 

5453 

5465 

5478 

5490 

5502 

5514 

5527 

5539 

5551 

2 

567 

9  10  11 

36 

5563 

5575 

5587 

5599 

5611 

5623 

5635 

5647 

5658 

5670 

2 

567 

8  10  11 

37 

5682 

5694 

5705 

5717 

5729 

5740 

5752 

5763 

5775 

5786 

2 

567 

8  9  10 

38 

5798 

5809 

5821 

5832 

5843 

5855 

5866 

5877 

5888 

5899 

2  3 

567 

8  9  10 

39 

5911 

5922 

5933 

5944 

5955 

5966 

5977 

5988 

5999 

6010 

2  3 

457 

8  9  10 

40 

6021 

6031 

6042 

6053 

6064 

6075 

6085 

6096 

6107 

6117 

123 

456 

8  9  10 

41 

6128 

6138 

6149 

6160 

6170 

6180 

6191 

6201 

6212 

6222 

123 

456 

789 

42 

6232 

6243 

6253 

6263 

6274 

6284 

6294 

6304 

6314 

6325 

123 

456 

789 

43 

6335 

6345 

6355 

6365 

6375 

6385 

6395 

6405 

6415 

6425 

123 

456 

789 

44 

6435 

6444 

6454 

6464 

6474 

6484 

6493 

6503 

6513 

6522 

1  2  3 

456 

789 

45 

6532 

6542 

6551 

6561 

6571 

6580 

6590 

6599 

6609 

6618 

2  3 

456 

7  8  9 

46 

6628 

6637 

6646 

6656 

6665 

6675 

6684 

6693 

6702 

6712 

2  3 

456 

778 

47 

6721 

6730 

6739 

6749 

6758 

6767 

6776 

6785 

6794 

6803 

2  3 

455 

678 

48 

6812 

6821 

6830 

6839 

6848 

6857 

6866 

6875 

6884 

6893 

2  3 

445 

678 

49 

6902 

6911 

6920 

6928 

6937 

6946 

6955 

6964 

6972 

6981 

2  3 

445 

678 

50 

6990 

6998 

7007 

7016 

7024 

7033 

7042 

7050 

7059 

7067 

2  3 

3    5 

678 

51 

7076 

7084 

7093 

7101 

7110 

7118 

7126 

7135 

7143 

7152 

2  3 

3    5 

678 

52 

7160 

7168 

7177 

7185 

7193 

7202 

7210 

7218 

7226 

7235 

122 

3    5 

677 

53 

7243 

7251 

7259 

7267 

7275 

7284 

7292 

7300 

7308 

7316 

122 

3    5 

667 

54 

7324 

7332 

7340 

7348 

7356 

7364 

7372 

7380 

7388 

7396 

122 

3    5 

667 

0 

1 

2 

3 

4 

5 

6 

7 

8 

9 

12  3 

456 

789 

Logarithms. 


Proportional  parts 


Nat. 
Kinbtr 

1° 

1 

2 

3 

4 

5 

6 

7 

8 

9 

1    2   3 

456 

r  s  9 

55 

7404 

7412 

7419 

7427 

7435 

7443 

7451 

7459 

7466 

7474 

1    2    2 

345 

5     6     7 

56 

7482 

7490 

7497 

7505 

7513 

7520 

7528 

7536 

7543 

7551 

122 

345 

567 

57 

7559 

7566. 

7574 

7582 

7589 

7597 

7604 

7612 

7619 

7627 

122 

345 

567 

58 

7634 

7642 

7649 

7657 

7664 

7672 

7679 

7686 

7694 

7701 

112 

3     4     4 

567 

59 

7709 

7716 

7723 

7731 

7738 

7745 

7752 

7760 

7767 

7774 

112 

344 

567 

60 

7782 

7789 

7796 

7803 

7810 

7818 

7825 

7832 

7839 

7846 

112 

3     4     4 

5     6     6 

61 

7853 

7860 

7868 

7875 

7882 

7889 

7896 

7903 

7910 

7917 

112 

344 

566 

62 

7924 

7931 

7938 

7945 

7952 

7959 

7966 

7973 

7980 

7987 

1     1    2 

3      3     4 

566 

63 

7993 

8000 

8007 

8014 

8021 

8028 

8035 

8041 

8048 

8055 

112 

334 

5     5     6 

64 

8062 

8069 

8075 

8082 

8089 

8096 

8102 

8109 

8116 

8122 

1    1    2 

334 

556 

65 

8129 

8136 

8142 

8149 

8156 

8162 

8169 

8176 

8182 

8189 

1    1    2 

334 

556 

66 

8195 

8202 

8209 

8215 

8222 

8228 

8235 

8241 

8248 

8254 

112 

334 

556 

67 

8261 

8267 

8274 

8280 

8287 

8293 

8299 

8306 

8312 

8319 

112 

334 

556 

68 

8325 

8331 

8338 

8344 

8351 

8357 

8363 

8370 

8376 

8382 

112 

334 

456 

69 

8388 

8395 

8401 

8407 

8414 

8420 

8426 

8432 

8439 

8445 

112 

234 

456 

70 

8451 

8457 

8463 

8470 

8476 

8482 

8488 

8494 

8500 

8506 

112 

234 

456 

71 

8513 

8519 

8525 

8531 

8537 

8543 

8549 

8555 

8561 

8567 

1    1    2 

234 

455 

72 

8573 

8579 

8585 

8591 

8597 

8603 

8609 

8615 

8621 

8627 

1    1    2 

234 

455 

73 

8633 

8639 

8645 

8651 

8657 

8663 

8669 

8675 

8681 

8686 

112 

234 

455 

74 

8692 

8698 

8704 

8710 

8716 

8722 

8727 

8733 

8739 

8745 

1    1    2 

234 

455 

75 

8751 

8756 

8762 

8768 

8774 

8779 

8785 

8791 

8797 

8802 

1    1    2 

2     3     3 

455 

76 

8808 

8814 

8820 

8825 

8831 

8837 

8842 

8848 

8854 

8859 

1    1    2 

233 

455 

77 

8865 

8871 

8876 

8882 

8887 

8893 

8899 

8904 

8910 

8915 

1    1    2 

233 

445 

78 

8921 

8927 

8932 

8938 

8943 

8949 

8954 

8960 

8965 

8971 

112 

2     3     3 

4           5 

79 

8976 

8982 

8987 

8993 

8998 

9004 

9009 

9015 

9020 

9025 

112 

233 

4           5 

80 

9031 

9036 

9042 

9047 

9053 

9058 

9063 

9069 

9074 

9079 

1     1    2 

233 

4            5 

81 

9085 

9090 

9096 

9101 

9106 

9112 

9117 

9122 

9128 

9133 

112 

233 

4           5 

82 

9138 

9143 

9149 

9154 

9159 

9165 

9170 

9175 

9180 

9186 

1    1    2 

233 

4           5 

83 

9191 

9196 

9201 

9206 

9212 

9217 

9222 

9227 

9232 

9238 

112 

233 

4           5 

84 

9243 

9248 

9253 

9258 

9263 

9269 

9274 

9279 

9284 

9289 

112 

233 

4           5 

85 

9294 

9299 

9304 

9309 

9315 

9320 

9325 

9330 

9335 

9340 

1    1    2 

233 

4           5 

86 

9345 

9350 

9355 

9360 

9365 

9370 

9375 

9380 

9335 

9390 

112 

233 

4           5 

87 

9395 

9400 

9405 

9410 

9415 

9420 

9425 

9430 

9435 

9440 

0    1 

223 

3 

88 

9445 

9450 

9455 

9460 

9465 

9469 

9474 

9479 

9484 

9489 

0    1 

2     2     3 

3 

89 

9494 

9499 

9504 

9509 

9513 

9518 

9523 

9528 

9533 

9538 

0     1 

223 

3 

90 

9542 

9547 

9552 

9557 

9562 

9566 

9571 

9576 

9581 

9586 

0     1 

223 

3 

91 

9590 

9595 

9600 

9605 

9609 

9614 

9619 

9624 

9628 

9633 

0     1 

223 

3 

92 

9638 

9643 

9647 

9652 

9657 

9661 

9666 

9671 

9675 

9680 

0    1 

223 

3 

93 

9685 

9689 

9694 

9699 

9703 

9708 

9713 

9717 

9722 

9727 

0    1 

223 

3 

94 

9731 

9736 

9741 

9745 

9750 

9754 

9759 

9763 

9768 

9773 

0    1 

2     2      3 

3 

95 

9777 

9782 

9786 

9791 

9795 

9800 

9805 

9809 

9814 

9818 

Oil 

223 

3 

96 

9823 

9827 

9832 

9836 

9841 

9845 

9850 

9854 

9859 

9863 

Oil 

223 

3 

97 

9868 

9872 

9877 

9881 

9886 

9890 

9894 

9899 

9903 

9908 

Oil 

223 

3 

98 

9912 

9917 

9921 

9926 

9930 

9934 

9939 

9943 

9948 

9952 

Oil 

223 

3 

99 

9956 

8961 

9965 

9969 

9974 

9978 

9983 

9987 

9991 

9996 

Oil 

2      2      3 

3      3 

- 

0 

1 

2 

3 

4 

5 

6 

7 

8 

9 

1   2  3 

456 

789 

157 


APPARATUS  AND   SUPPLIES 


THE  estimate  appended  has  been  kindly  furnished,  at 
the  author's  request,  by  George  D.  Feidt  &  Co.,  528  Arch 
Street,  Philadelphia. 

1.  Apparatus  for  each  student  (cost,  $5.95). 
Bunsen  burner,  test-tube  brush,  retort  stand   (small, 

two  rings),  clamp,  deflagrating  spoon  (small  iron),  file 
(triangular),  blow-pipe  (brass,  common),  file  (round), 
iron  wire  gauze,  5"  X  5",  tin  pan,  round,  10"  X  3"  high, 
lead  disc  for  collecting  gases,  2J"  diam.,  4  lead  rings  to  fit 
6  X  t"  test-tube,  100  c.c.  evaporating  dish,  mortar  and 
pestle  (3"),  2  beakers  (straight  form,  no  lip,  100  c.c. 
capacity),  funnel  3",  gas  generating  bottle,  capac.  250  c.c., 
with  funnel  tube  and  delivery  tube,  £  doz.  plates  of  win- 
dow glass  4"  X  4",  4  salt-mouth  bottles  300  c.c.  (common, 
no  stoppers),  J  doz.  test-tubes  6  X  f ,  i  doz.  test-tubes 
8  X  1,  1  test-tube,  hard  glass,  6  X  i,  1  test-tube  rack  to 
fit  6  X  f  and  8X1  tubes.  Set  4  bottles,  1  liter  capacity, 
enameled  or  ground-glass  labels,  for  hydrochloric  acid, 
sulphuric  acid,  nitric  acid,  and  ammonia.  Wooden  block 
for  holding  these  bottles,  four  round  depressions  in  block. 

2.  Apparatus  to  be  kept  in  stock  anMfcjnished  as 
occasion  requires. 

A.  METAL   (cost,  $9.78,  the  induction  coil  being  the 
chief  item). — |  doz.  Hoffman  screws,  lead  dish  (flat,  75 
c.c.),  small  horseshoe  magnet,  wire  cutter,  induction  coil 
£"  spark). 

B.  PORCELAIN  (cost,  $7.20). — Porcelain  crucible  and 
cover  No.  0,  3  cells  (Edison-Lalande  or  equivalent). 

159 


160 


ELEMENTARY  CHEMISTRY 


C.  GLASS  (cost,  $16.16).— Beaker   (500  c.c.),  beaker 
1,000  c.c.,  bulb  tube  (1"  bulb),  burette  50  c.c.  grad.  in 
•jij-  c.c.,  Hoffman  apparatus  for  electrolysis,  lamp  chimney 
(ordinary  wide  form),   U-shaped  eudiometer,  flask  (1,000 
c.c.),  flask  250  c.c.,  cylinder  graduated  1,000  c.c.,  Erle- 
meyer  flask  1,000  c.c.,  separating  funnel  100  c.c.,  retort, 
capacity  200  c.c.  with  tubule  and  glass  stopper,  U-tube  4" 
with   side   tubes,    glass    rod,    glass    tubing.     Eudiometer 
straight  (50  c.c.  graduated  in -J- c.c.). 

D.  VARIOUS    (cost,    $7.20). — Corks,    wooden    (best), 
corks,  rubber,  1  hole,  2  hole,  and  solid,  asbestos  board, 
Christmas-tree  candles,  filters  (common  5"),  labels  (Den- 
nison's  No.  209),  magnifying  lens,  meter  stick,  yard  stick, 
platinum  foil,  platinum  wire   (size  for  blow-pipe  work), 
copper  wire  (insulated),  iron  wire,  rubber  tubing  J"  inter- 
nal diam.,  ^"  internal  diam.,  \"  internal  diam.,  and  large 
for  gas  connections. 

CHEMICALS  (cost,  $23.70) 

Acid,  acetic, 
"     hydrochloric, 
"     nitric, 

"     sulphuric  (all  of  above  commercial), 
"     hydrochloric, 
nitric, 

sulphuric  (all  of  above  G.  P.), 
hydrofluoric,  commercial, 
oxalic, 
tartaric. 

Alcohol,  95  per  cent, 
"       wood  spirit. 
Aluminium,  sheet, 

powder,  bronze  paint, 
sulphate. 

Ammonium,  chloride, 
nitrate, 

"  sulphate, 

"  oxalate. 


APPARATUS  AND  SUPPLIES  161 

Ammonia  water. 
Aniline  red. 
Antimony,  metallic, 

potass,  tartrate  (tartar  emetic). 
Arsenic,  metallic, 

"        trioxide. 
Barium  chloride, 
chlorate, 
"        peroxide. 
Bleaching  powder. 
Bone-black. 
Borax. 
Bromine. 

Cadmium  chloride. 
Calcium  carbide, 

"        carbonate  (cracked  marble), 

"        chloride  (crystals), 

"        chloride  (fused), 

"        fluoride  powder, 

"        oxide  (lime), 

"        sulphate  (plaster  of  Paris). 
Cannel  coal. 
Carbon  disulphide. 
Charcoal,  lump. 
Chrome  alum. 
Cobalt  nitrate. 
Cochineal. 
Copper,  clippings, 

"        sheet, 
wire, 

"        carbonate,  basic  (powder); 

"        sulphate. 
Cream  of  tartar. 
Ether. 
Gold  leaf. 

Hydrogen  peroxide  (3  per  cent). 
Iodine. 
Iron  (alcoholized). 

Steel  wool. 
Iron  chloride  (ferric), 
"    oxide  (caput  mortuum), 
"    oxalate  (ferrous), 


162  ELEMENTARY  CHEMISTRY 

Iron  filings, 

"    sulphide. 
Kerosene. 
Lead  (tea  lead), 
"    granulated, 
"    acetate, 
«    red  lead, 
"    dioxide, 
"    nitrate, 
"    oxide  (litharge). 
Litmus  (cube). 
Litmus  paper,  red  and  blue. 
Magnesium  oxide, 
"          powder, 
"  ribbon. 

Manganese  dioxide  (powder), 

"  "        (granulated). 

Mercuric  chloride, 
"          iodide, 
"          oxide. 
Mercurous  nitrate. 
Phenol-phthalein. 
Potassium  (metal), 
"          bromide, 
"          carbonate, 
chlorate, 
chloride, 
"          chromate, 
"          dichromate, 
"          ferricyanide, 
"          hydroxide, 
"          iodide, 

permanganate, 
sulphate, 
sulphocyanide. 
Pyrogallol. 
Rosin. 

Silver  nitrate. 
Sodium  (metallic), 
"       acetate, 

"       dicarbonate  (baking  soda), 
"       carbonate  (crystals), 


APPARATUS  AND  SUPPLIES  163 

Sodium  carbonated"  (dry), 
"        chlorate, 

chloride  (salt), 
"        hydroxide, 
"        nitrate, 
"       nitrite, 
"       silicate, 
"        sulphate, 
"        acid  sulphite, 
"        thiosulphate. 
Starch. 

Strontium  chlorate. 
Sugar,  grape  (glucose). 
Sulphur,  roll. 

"         flowers, 
Tin,  granulated, 
"     foil, 
"     bar. 

Turmeric  paper. 
Vaseline. 
Water,  distilled. 
Zinc,  dust, 

"     granulated  (mossy), 
"      sheet, 
"     oxide. 

All  prices  are  net.  The  prices  of  the  chemicals  are 
subject  to  variation.  The  quantities  of  the  chemicals  will 
be  sufficient  for  a  class  of  six  or  more  to  make  the  experi- 
ments. 


(5) 


rcu 


D  LD 


APR  2  7 


YB   16923 


THE  UNIVERSITY  OF  CALIFORNIA  LIBRARY 


